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OSMANIA UNIVERSITY LIBRARY
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NEW WORLD OF
CHEMISTRY
SCIENCE IN THE SERVICE OF MAN
EW
We
ORLD OF
HEMISTRY
BERNARD JAFFE
SILVER BURDETT COMPANY
NEW YORK - DALLAS • CHICAGO • SAN FRANCISCO
BERNARD JAFFE
Chairman, Drpmtmrnl of j'hyuftil Science
Jnmr\ Wutli\nn ////,'/< .SV hoitl, AV/i York (Iity
AUTHOR OF
MEN OF SCIENCE IN AMERICA
OU iPOSTS CF SCIENCE
CRUCIBLES
CHEMICAL CALCULATIONS
Copt rif fir. l<>t?, JO tO. 107.?. /077
SILVER BURDETT COMPANY
[\K\\ \\om.l) OF r.llhMISTRYua* In-r pulili^linl in lU3i
and ri»ni|ilrtrh ir\i->rd in L(MJ and 11U7. Minor icvi^ion-
lia\«- IMTII madr r\n\ t\\o or llirrc \«MI- to krr|» (lu* hook
up-(o-dalr 'l'ln-« I1).").") nlition i>« .in fxtcn^ue rr\iNioii 'I'ht- tr\t
IKKS IM>CII Kirp-l\ rruiitlcn, nr\vl\ lio^i^ncd. nculx illn-lratrd.
Format Design . . . OliMe Wlntnrv
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Printed by lMiin|iton l'rr*s
I
O THE TEACHER:
. For success and satisf action in modern living, tin* citi/en
must understand and appreciate the scientific aspects of his
environment and the role ol science in the development of
civili/ation. lie must hase his everydav thinking and action
on the best information available with a full reali/ation of
why the methods of science are superior to other methods of
obtaining information. Finally, the citi/en must be ready to
use scientific knowledge lor the <M)od of all. M \v \\'ORI,|) or
CIII.MISTKY is designed to help students achieve these objec-
tives. It is a textbook for yoimjj; people who are learning to be
citi/ens.
This 1 ()r).") edition of MAV \VOKI.D 01- CIIMMISTRY lias been
brought abreast of recent chemical developments and dis-
coveries. Striking advances have been made in nuclear energy,
petrochemistry, metallurgy, textiles, and plastics. The chapters
dealing with these subjects have been tlionni^lHy revised to
take these changes into account.
The basic plan ol the book, which has been so widely
accepted, lias not been altered. Hut within that framework
many changes have been made. Diagrams have been com-
pletely redrawn and enlarged. New illustrations have been
provided. The index lias been expanded to make the text an
even more effective source book. Despite the addition of new
materials, however, the hook as a whole lias been shortened.
Perhaps most important of all. the chapters dealing with
the basic theory and mathematics of chemistry have been
rewritten. Formulas, valence, equations, and problems have
been £iven a new and expanded treatment. The chapters pre-
senting the electron theory, ioni/ation. and the study of teases
have been rewritten to achieve a maximum of clarity. While
the concepts of chemistry themselves cannot be simplified.
vi TO THE TEACHER
thry can he made r.isiri to understand h\ hcttcT arrangrmrnt
and tfn.ilci clanlx of laniiua^c' Glarit) of wilting has been
.L m.ijoi 140.1! of this substantial levisinn of NKW WORLD OF
CllhMIVlRY.
'I he boc»k has been \\iitten \\ith the1 puipose of getting idea*
acrms to the \tudcnL With the exception of ncccssaix teehnical
trims, tin \o(,ibulai\ consists of familial words. The style is
duett, tin sentences and paia<;iaphs are short. Much of the
illiistiatm mateiial has been draxxn fiom the student's own
enviionmc nt (lonsumii aspects of che-mistix and the them-
isti\ of common things aie stiessed at the same time that
ic(|iiiienie.its of model n couises of stud\ and examining boards
aie iulh met.
Select! d matenals from the lnstor\ of c hemistr\ arouse
inteiest .ind enable the student to sec how fundamental scien-
tific ideas ha\e developed and expanded. Thc\ revc\il that
rliemic .il thioiies ,md principles arc* descriptive generali/ations
of ni. ill's oinv ing i xpiMieni i in tr\mi> to undei stand his world.
FinalK, the1 hi^tonc materials hel]) studrnts scu- hem scientific
discoxm has aflerte-d modem civili/ation and hem the1 needs
of sot iet\ c'onstantK stimulate* s( lentific rest-. in h
The* authoi \\islics to expiess his thanks to the mail}
teachcTs. diemisK and industrial, e'diu ation.il and govc-rn-
menttil or^.un/ations that have1 generousU made suggc-stions,
chicked the acciuacA of mateiials. and provided illustiations
and data foi this nc\\ edition. He aNo \\ishes to ac'kncm ledge*
his indebtedness to his son. I)i Lionel F Jaffe-, and to Lrc
I)ei»hton and John H \Villi.iniscm of Silver Buidett (Company
\\liosi- i-aieliil plaiinui!*. sound judgmc^nl, and nrvcr-f ailing
rntliiisia^ni h.ixe made this c-oopeiative- efTort e-xtie-me^h pleas-
ant and, it is hoped, eflrctive
Ft nun id Jaffe
CONTENTS
1. MATTER and Its Changes I
What the \\oild is made of .uul ho\\ M icntilu methods expand
our undci^tandmg.
2. OXYGEN: Earth's Most Abundant Element 24
Distovei), piopeities, and uses of the gas \\lmli is the bicath of
life.
3. HYDROGEN: Lightest of the Elements 44
The piimat\ stnlT of \\hkh tlie entne uni\eise is l)inlt.
4. WATER: Most Common Liquid Compound 60
The thiillmg discoveiy of its composition, its ptopeitics, and uses.
5. ATOMS: Bricks of the Universe 75
\ Quaker sc hoolteat her explains t\\o laus of <hemisti\ b\ means
of his atoms
6. FORMULAS: The Chemist's Abbreviations 85
Rules on the untin^ of chemical foimulas.
7. ATMOSPHERE: The Ocean of Air 97
I low the components of the air \\ere discovered and utili/ed.
8/ EQUATIONS: Shorthand of Chemistry 114
Rules and aids in the halancing of chemical equations
9. MATHEMATICS of Chemistry 124
Meaning of atomic wriijhts and how some niathrm.itic.il pioblems
aie solved.
10. CHLORINE and the Halogen Family 137
A closelv related group of elements — fluorine, c hloi me, bromine,
and iodine.
I L ELECTRONS and Other Particles 154
Origins, development, and chemical usefulness of the latest theory
of the structure of the atom.
viii CONTENTS
12. NUCLEAR ENERGY at Last! 176
M.m hn.ilU p<netiates the < enter ot the atom and produces nu< lear
changes
13. ACIDS: Hydrochloric Acid, a Typical Acid 195
PiopeitKs .iiid g<nei,d method of piep.ii m«j a< ids UK hiding IIF
14. BASES: Sodium Hydroxide, a Typical Base 206
Nrutr ah/at ion of bases b\ .Kids, and salt formation
15. SOLUTIONS: Water, the Universal Solvent 217
Kinds of solutions, distilled and hid:\ \\atei
16. IONS and Dissociation 231
\\n\\ disputation explains stieimth ol acids, ludiol\Ms, electrolysis,
17. AMMONIA and Reversible Reactions 250
I lie lommeiu.d s\nthesis of aninioma a milestone in <hemistr\
18. NITRIC ACID and Nitrogen Compounds 262
1 he (omiiKKial s\nthesis of nitiit .Kid, and mttogeii fixation
19. MOLECULES: Avogadro's Hypothesis 276
Imliidmu applications lo the solution of \\ni;lit-\oltinir and
sliai^hl-\ olume pioblenis
20. SULFUR and Hydrogen Sulfide 289
I hen piodmtion, piopeities, and uses, .IK hiding the manuf.ieture
of mat( lies
21. SULFURIC ACID: The Fundamental Acid 304
Its (ommeui.il piepaiation, piopeities, and mam uses.
22. ALLOTROPIC CARBON: Key Source of Energy 320
Piopeities and uses \\ith spei i.il attention to coal
23. CARBON DIOXIDE: Gas of Life and Decay 339
Including caibonic aeid, its salts, and baking poudeis.
24. CARBON MONOXIDE and Other Gaseous Fuels 352
Including water gas, producer gas. coal gas, natuial gas, acctvlenr
25. METALS and Their Chemical Activity 369
Including the piepaiation, piopcities, uses, and detettion In the
spectroseope of sodium, potassium, .ind lithium.
CONTENTS ix
26. ALUMINUM: Most Common ol Light Metals 387
I IK llldllli! MIMIC <>t Its compounds, .Hid j|s0 {\\c lUCl.il Ill.t
27. IRON and Steel 404
Intituling .1 ulmipse of the ne\\ linn/mis in the steel mdiis(i\.
28. COPPER: Nerves of the Machine Age 423
Its met. il hit <j,\ and uses, iiu hiding coppei siill.itr.
29. OTHER METALS and Their Uses 436
Including the nxe of the M>-<.ilh(i iaie metals.
30. FERTILIZERS and Salts of Sodium 464
IIK lulling the solution ol oin pot.ish piohlcin.
31. CALCIUM: Its Common Compounds 480
u supc iphiisph itcs .UK! \\.itrr
32. IRON: Some Special Compounds 496
I IK ludinu then uses in inks .ind l)luc|)iint ]).»per.
33. GLASS and Some Silicon Compounds 505
IIH liiciinu ,dso some hoTon ( on i pounds MM h .is borax.
34. HYDROCARBONS and Their Derivatives 522
n^ the s|)C( t.n nl.ii iis<- of petioi heinistix .1 new
35. ALCOHOL and Other Organic Compounds 547
Anothei i*limp\r of th<- \\cnlcl of the OII^.IIIM < hnnist.
36. FOODS and Chemotherapy 566
Chemistry .it \\oik in the ser\ue of m,m\ he.dth
37. FIBERS AND PLASTICS: Textiles, Paper, and Dyeing 591
Including the ne\\ \\oild of in.iii-ni.idt
38. COLLOIDS: The Colloidal State of Matter 610
The rhcnustiy of tiny paitic Irs drops. fiLiincnts, L^i.uns, and films
39. LIGHT: Its Chemical Effects 623
Including the rssrnti.il principles of sunple photnt^i.tphy
40. MORE CALCULATIONS 636
Simplest formula and true formula di trniimmc; atomic s\eit^ht ex-
perimentally, temperature s<ale <on\ersjon, use of Hoyle's Jaw and
(ihaile^ la\>.
MATTER:
AND ITS CHANGES
. . . We must trust in nothing but
facts. These are presented to us by
nature and cannot deceive. We ought
in every instance to submit our
reasoning to the test of experiment.
It is especially necessary to guard
against the extravagances of imagina-
tion which incline to step beyond
the bounds of truth. Antoine Lau-
rent Lavoisier, 1743-1794
Progress through scientific knowledge. A play, presented in New
York City, is seen and heard instantly and simultaneously in millions
of homes throughout the nation, as clearly as if its vast audience
were actually seated in the studio. An airplane takes off from an
airfield in California and within a few minutes is screaming through
the stratosphere 60 thousand feet above the earth at more than 700
miles per hour. In an Alabama farmhouse, a country doctor diag-
noses his patient's illness as pneumonia, and assures the family that
it is quickly curable by the administration of certain drugs. In an
isolated group of islands in the South Pacific, a bomb explodes with
a force equal to hundreds of thousands of tons of dynamite. In Iowa,
a housewife cooks a salt-water fish caught by her husband during the
family vacation six months before and two thousand miles away.
A generation ago, any of these occurrences would have been con-
sidered miraculous. Yet today they arouse no unusual public excite-
ment, for they are no longer extraordinary events. To us they are
familiar happenings, representative of what we call the progress of
ivilization.
Progress has been taking place in some measure ever since man
appeared on earth. But within the past two or three centuries,
1
NEW WORLD OF CHEMISTRY
civilization has progressed much more rapidly than in all the previ-
ous thousands of years of man's history, with the most radical changes
occurring within the past few decades. In the years ahead, civilization
will continue to progress. More of the "incurable" diseases will be
conquered; methods of transportation and communication will be
further improved; the force which gives the atomic and hydrogen
bombs their great destructive power will be harnessed for man's
welfare; our everyday lives will be made safer and more com-
fortable.
What enables us to perform acts today which were unheard of
just a few short years ago? How may we be certain that the future
will bring forth new wonders? The answer to both of these questions
lies in the knowledge we now possess concerning the nature and
behavior of those things which make up the entire world and all its
living creatures. In short, the answer lies in our scientific knowledge.
Scientific knowledge is not something which has been created
recently. It has been gathered by many men, known and unknown,
during all of man's centuries on earth. Just as a snowball rolling
downhill is at first small and slow moving, but in time increases both
its si/e and speed, so it has been with scientific knowledge. In his
early days on earth, man knew a limited number of scientific facts,
increasing the number with painful slowness as the centuries went by.
But gradually, almost imperceptibly, the number of facts became
larger and their rate of discovery quickened. Within the past few
hundred years, particularly during the twentieth century, they have
been acquired at a prodigious rate. New scientific advances are an-
nounced almost daily. A rolling snowball must finally come to rest
at the foot of a hill, but there is no indication that scientific knowl-
edge will ever cease to grow. Nor does there appear to be any limit
to the progress civilization can make through the application of this
knowledge.
Linde Air Products Company
(left) The highly-trained
professional chemist makes
many important contribu-
tions to civilization.
(right) The lessons of the
student chemical labora-
tory find many useful ap-
plications in daily life.
MATTER AND ITS CHANGES 3
Chemistry has helped make a new world. No single science is
entirely responsible for our modern civilization. But chemistry ranks
high among those which are considered the most important because
it is a basic science^ essential to virtually every scientific study regard-
less of its nature. In the manufacture of every product of our great
industrial civilization, chemistry plays ajvital role.
The work of the chemist affects all of us — the physician, the
dentist, the engineer, the public health employee, the soldier, the
gas station attendant, the clerk, the housewife, the worker of factory,
jninc, or farm, the schoolboy and the schoolgirl.
While the achievements of modern chemists seem at times to be
the result of some kind of magic, nothing could be further from the
truth. Chemical ^magic'^is the result of years of study, hard work.
and struggle; of burning the midnight oil when it seems a problem
has no solution; and, finally, of a thrilling moment when everything
hinges on one more experiment and the world spins 'round in a test
tube or flask. Nor are all experiments rewarded with striking suc-
cess. Sometimes they fail and the chemist must begin, wearily, but
with determination, to retrace his weeks or months ol work in an
attempt to discover where he went astray.
Chemistry is an intensely interesting and rapidly changing field
and, like all sciences, its roots extend far back into history. Today's
chemist owes a great debt to the men who, over many centuries,
added to the knowledge he now uses in performing his "miracles."
The lives of those men, how they accumulated information and how
that information is applied in modern chemistry is a thrilling story,
filled with suspense and drama. It is the story told in NKW WORLD 01-
CHEMISTRY.
Although we live in the "scientific age," we still see, everywhere
in the world, superstitions and prejudices, hunger and disease. To
correct these conditions, we must use in the outside world what we
learn in the classroom. We must base our actions upon scientific
knowledge, always bearing in mind, however, that we know little
with absolute finality.
Board of Education, City of New York
4 NEW WORLD OF CHEMISTRY
We must also remember that science is a two-edged sword which
may be used either to serve man or to destroy both him and his
^ivbrks. Although the spirit of science is essentially democratic and
constructive, we alone can prevent its becoming an oppressively
tyrannical and horribly destructive weapon. Not many of us may
ever become professional scientists, the men and women who work
in the great laboratories of industry, government, and education.
But we can all become scientists in an even broader sense.
We can act on the most dependable information available, using
the searching light of science to wipe out prejudices, half truths,
and incorrect beliefs. The methods and knowledge of science are in
our hands. We can live more fully, more satisfyingly, more com-
fortably, more humanely, and more intelligently only by using them.
Look forward to the future, with the assurance that, through the
j^fforts of each of us, science amafld w ill serve mahkTncT
What is the world made of? For thousands of years, men have
tried to find out what the material things of the world are made of.
Socrates (sok'rd-tez) , a learned Greek teacher who lived in the fifth
century B.C., believed that he might discover the answer to this ques-
tion simply by thinking about it. One other Greek teacher is said
to have put out his eyes so that his thinking might not be disturbed
or influenced by what he could see about him. He wanted to give
himself over completely to undisturbed thinking, or contemplation.
To be sure these were extreme cases. Not all the ancient teachers
relied wholly on pure thought as they tried to find out what the
world is made of. Some of them, of course, were influenced by what
they could observe. For example, Thales (tha'lez) , who was born in
Miletus in Asia Minor in 624 B.C., noticed that water nourishes crops
and that it is found in large amounts in the bodies of men and other
living things. Hence he speculated, or guessed, that water was the
basic substance from which all material things were made.
The speculations of Thales were more important to him than his
observations. It is said that one day, as he was walking along ab-
sorbed in deep thought and looking up at the sky, he fell into a well.
Thratte, a housemaid, saw the accident and laughingly said: "In his
zeal for things in the sky, he does not see what is at his feet." No
doubt, most of you know how easy it is to become so absorbed in
trying to solve a problem that you become "blind" to common, every-
day things and may even be called "absent-minded." In a similar
manner, it is possible to become so engrossed in thinking about a
problem that simple methods of solving it do not occur to the
thinker. Some of the ancient teachers were victims of this "blindness."
Careful experiments and
constant observation are
essential to the chemist's
study of matter.
Many of Thales' conclusions concerning what the world is made
of were wrong. He and other scholars like him depended too much
upon what is now called abstract thought, and altogether too little
upon careful experimentation and observation. II they had depended
less upon abstract thought and more upon observation and experi-
mentation, accurate information would have accumulated much
more rapidly. However, even though many of their conclusions were
wrong, their work was of value. They were thinking about the things
that surrounded them, which was more than was being done by most
of the people who lived at that time.
Of course, in answering any question, careful thinking is neces-
sary, but it must be thinking based_iipon act 'uralc Jacls, which can
comc^only from careful observation. Many of the ancient teachers
were fine thinkers. They began with what they knew, or thought
they knew, and reasoned logically to a conclusion. Many of their
conclusions were false, however, because they based their thinking
on inaccurate or incomplete information.
Ours is a complex world. This age-old search to find out what the
world is made of is not over. Far from it. You, living 2500 years after
Thales, know that water is not the basic substance from which the
world is made. Yet, if you were asked to name all the different ma-
terials that make up the world, you would doubtless think you had
been given an impossible task. Imagine listing all the materials that
make up the rocks and soil beneath your feet, the vast expanse of
water that covers three-quarters of the earth, the multitude of liv-
ing things, the deep envelope of air that surrounds our globe, and
the billions of stars, some so far from the earth that their distances
stagger our imagination!
At extremely low tempera-
tures, a gas may become
liquid or even solid. This
flask contains liquid helium
and solid air. Because the
frozen air at a temperature
of —340° F. is much "hot-
ter" than the liquid helium
at -452° F., it is causing
the helium to boil and over-
flow the flask.
Westinghouse Electric Corporation
Matter exists in three states. Even though you could not name all
the substances that make up our world, you would at least know that
some of them exist as solids, others as liquids, and still others as
gases. In fact, probably you know that mailer, which is anything that
has weight and takes up space, occurs in one of three conditions —
solid, liquid, or gaseous. In which of these conditions, or slates of
matter, any substance exists depends partly on the nature of the sub-
stance itself, partly on its temperature, and partly on the pressure.
Many substances exist in all three states, depending on the tem-
perature. For example, water (a liquid) may be changed to ice (a
solid) by cooling, or to water vapor (a gas) by heating. Changes
from one state of matter to another by heating or cooling are very
common. Iron, which we know as a hard gray solid, is melted in
foundries and changed to a shimmering, silvery liquid. If its tem-
perature is raised high enough, gaseous iron vapor is boiled oft.
While iron vapor is something with which most of us are not fa-
miliar, astronomers know that iron exists normally in the gaseous
state in certain extremely hot stars.
What is a physical change? When water changes to steam or to
ice, only its form has been changed. Steam is a form of water; ice is
also a form of water. When a piece of limestone is pulverized, only
its form has been changed. The small pieces are still limestone, al-
though their form differs from that of the huge chunks in which they
came from the quarry. These changes in the water and in the lime-
stone have been in form only. Neither the water nor the limestone
MATTER AND ITS CHANGES 7
was changed into another substance. A change in which the original
substance does not change into one or more other substances is a
physical change.
Break a piece of wood in two. Heat a piece of iron wire in a vac-
uum by passing an electric current through it. The heat produced
causes the wire to glow, but, after the current is shut off, the wire
returns to its original condition. In each case, what kind of change
has occurred?
How does a chemical change differ from a physical change? You
know that when a piece of paper burns, it is completely and radi-
cally changed. The hot gases that are given off and the ash that is
left behind do not in any way resemble the original substance. When
gasoline is burned in an engine, the resulting substances are entirely
different from the liquid gasoline. Animal tissue is totally different
from the vegetable substances from which it is made. The dull tar-
nish on silverware differs completely from the gleaming silver.
In these changes more than the form of the original substance
was altered. In each case the composition of the original substance
changed. Some form of energy, usually heat, was either liberated
or absorbed. A change in which the original substance disappears
(changes] and new substances are formed is a chemical change.
Place a small quantity of sugar in a beaker or other glass vessel.
On the sugar, pour a small quantity of sulfunc acid (a liquid which
is discussed in detail in Chapter 13) . The white sugar changes to a
black spongy substance, which cannot be dissolved, or absorbed, in
water. What kind of change has occurred?
Chemistry is the science that deals with the composition of matter
and with the many chemical changes which matter undergoes.
How is a substance identified? Telling one substance from an-
other is called identifying a substance. When you try to identify a
substance or to find out whether a substance has undergone a chemi-
cal or physical change, you need to know its characteristics — what
the substance is like, and how it acts with other substances.
We find out what a substance is like by asking such questions as:
What is its color? its odor? its taste? Is it a solid? a liquid? a gas?
At what temperature does it boil? At what temperature does it
freeze? How hard is it? Does it conduct electricity? The answers to
these questions are characteristics that enable us to describe and
identify a substance. These characteristics of a substance are called
its physical properties.
We find out how a substance acts with other substances by placing
it in contact with these substances and observing what occurs. We
* NEW WORLD OF CHEMISTRY
ask also how sunlight, electricity, and heat affect it. Our observations
and the answers to these and other questions give additional charac-
teristics that enable us to describe and identify a substance. These
characteristics are called the chemical properties of that substance.
Since two different substances never have exactly the same physical
and chemical properties, any substance may be identified by deter-
mining these properties.
The ancients believed the world made of "four elements." But
to return to our original question: What is the world made of?
Guesses and speculations would be useless in attempting to answer
this question. Because the ancients depended chiefly upon these pro-
cedures and also upon inaccurate and uncontrolled experimentation
and observation, they made little progress in answering the question.
After Thales had suggested water, another man proposed that air
might be another of the basic substances from which all matter was
made. Fire, too, was suggested and later earth. Pythagoras (pi-thag'6-
ras) , an ancient Greek thinker and mathematician who lived about
600 B.C., is thought to have been the first European to express the
idea that all matter was composed of these "four elements."
These conclusions seemed to be proved by the observations of the
early investigators. When a stick of green wood was burned, they
saw that lire was produced, water was forced out and boiled off at
the ends of the stick, a smoky vapor (air) was given off, and an ash
(earth) was left behind. They concluded, therefore, that all matter
was made up of different amounts of two or more of these four basic,
or elementary, substances.
The Greek thinkers, however, made a serious mistake. They failed
to make enough observations of different substances. They did not
make enough experiments. Consequently, their conclusions were
wrong. Strangely enough, the idea that all matter is composed of
"four elements" (earth, air, fire, and water) persisted until the
eighteenth century and was considered correct by many otherwise
well-informed persons. Even today speakers and writers often refer
to the violent actions of air and water as "the fury of the elements."
The universe is made of an even 100 chemical elements. Scien-
tists now consider that the mountains, the oceans, the air, all living
things, and even the stars and the rest of the universe are composed
of simple natural substances that cannot be broken down, or de-
composed, into simpler substances by the ordinary types of chemical
change. A substance that cannot be broken down, or decomposed,
into a simpler substance by the ordinary types of chemical change is
an element.
These natural elements do not all occur on earth in equal
amounts. Taken together, 20 of them make up 99.5 percent of the
weight of the crust of the earth. All the other elements comprise
only O.f> percent of its weight. In all. there are 100 elements. Eight
of these elements have been produced in laboratories by scientists.
Probably the elements with which you are most familiar are gold,
silver, iron, copper, nickel, tin, aluminum, sulfur, oxygen, carbon,
nitrogen, and hydrogen. A list of all the elements is given on page
678.
What is a compound? Most of the substances we see, such as sand,
chalk, cotton, table salt, and water, are not elements. Rather, each
is composed of two or more elements so combined that (1) only
chemical action can tear them apart, and (2) the elements of which
each substance is composed can no longer be identified by their
original individual properties. A substance composed of two or more
elements .\o combined that the elements can no longer be identified
by their original individual properties is a compound. The elements
of which a compound is composed are said to be chemically com-
bined, or themically united.
Marble, for example, is a compound made up of three elements,
carbon, calcium, and oxygen, chemically combined. The properties
of a compound, such as color, odor, taste, form, and ability to dis-
solve in water, are nearly always distinctly different from the prop
erties of the elements of which it is composed. For example, pure
cane sugar, a sweet, white, crystalline solid which dissolves in water,
is completely different from any and all of the three elements of
which it is composed.
How does a mixture differ from a compound? In a compound
the elements must be chemically united. But there are other kinds
of substances made up of two or more elements or compounds. Al-
or mixed, each of the original substances can still be identified by
its original individual properties. Hence, the substances are not
chemically united.
A pinch of salt and a pinch of white sand stirred together make an
excellent example <>!' one of these substances. The salt can be iden-
tified by its characteristic taste and the sand by its gritty feel on the
tongue and teeth. A substance composed of two or more elements or
compounds that still retain their individual properties after they
have been thoroughly mixed is a mixture. Some1 of the most useful
substances in the world, such as soil, air, petroleum, and milk and
many other foods, are mixtures.
The properties of a mixture are the same as the properties of the
elements or compounds that compose it. A handful of iron powder
mixed with a handful of powdered sulfur makes a mixture that re-
sembles both the black iron and the yellow sulfur. If a magnet is
passed through it, the iron clings to the magnet. If a liquid (ailed
carbon disulfide is added to the mixture, the sulfur is dissolved. But
if the mixture of sulfur and iron is heated, these two elements com-
bine, forming a compound known as iron sulfide. Iron sulfide does
not look like either sulfur or iron. It is not magnetic and does not
dissolve in carbon disulfide. The properties of this compound do not
resemble those of either sulfur or iron.
Substances in certain mixtures may be separated mechanically. A
mixture of salt and sand may be separated by adding water. The
salt dissolves and the sand settles to the bottom. A mixture of iron
and sulfur may be separated by passing a magnet through it or by
adding carbon disulfide which dissolves the sulfur.
The phlogiston theory, an erroneous explanation ol burning.
One of man's greatest early achievements was the discovery of the
use of fire. So strange did fire appear that for a long time men wor-
shiped it. They considered it the force responsible for all creation.
They pondered over its mystery and made many attempts to explain
it. Karly alchemists thought that fire was the result of some vague
"sulfur" which burnable substances contain. But later alchemists
felt the need for a better explanation — an explanation which took
into account more of the facts that had been observed in burning
^ many different substances. A statement that takes into account and
attempts to explain observed tacts is known as a theory.
10
Over many centuries, alchemists, the forerunners of modern chemists, worked in
vain with their crude equipment to find the secrets of prolonging life and of
making gold from base metal.
About 300 years ago Becher (bek'er) , a German scientist, ad-
vanced the theory that all burnable substances contain phlogiston
(flo-jis't6n) , or "fire stuff." He said that when a substance burned,
phlogiston left it in the form of (lame. Becher thought that the ash
formed when a substance burned was the substance in in us its phlo-
giston. According to his theory, substances that burn readily, leaving
little ash. contain a great deal of phlogiston, while substances that
burn with difficulty and leave much ash contain little. The phlo-
giston theory was the first great theory in chemistry.
The phlogiston theory seemed correct to the alchemists because of
certain observations they had made. A rising candle flame seems to
tug at the wick. To the alchemists this suggested that phlogiston was
escaping from the binning candle. \Vhen a small amount of pow-
dered lead is heated in an iron spoon, it melts, burns, and forms a
yellow powder. According to the phlogiston theory, this yellow pow-
der is lead ash, or lead minus its phlogiston. Now if some way could
be found to add phlogiston to this lead ash, lead should be produced
again. Perhaps this could be done by heating the lead ash on some
substance that contains a lot of phlogiston, such as charcoal. The
charcoal might give up some of its phlogiston to the lead ash. When
this experiment is performed, the final product is actually lead.
You can see that this experiment seems to prove the correctness
of the phlogiston theory. Why? What mistake did the early sci-
entists make in attempting to prove the phlogiston theory?
For more than two centuries the phlogiston theory was considered
to be an accurate explanation of burning, and many of the famous
pioneers of modern chemistry, among them Priestley (prest'll) and
Schcele (sha'l<?) were its ardent supporters.
The first clue to the true explanation of burning. Kven though
<the phlogiston theory seemed to be upheld by experiments similar
to the one with lead, other experiments showed it to be false. In
these experiments the fact that many substances increase in weight
when burned could not be explained on the basis of the phlogiston
theory, for according to it, all substances lose phlogiston when
burned, and thereby lose weight.
11
12
Lavoisier Priestley
In 1774 Joseph Priestley, an English minister and amateur scien-
tist, led the way toward a true explanation of burning by his dis-
covery of a gas later named oxygen. Because of political and reli-
gious persecution, he fled from England and spent his last years in
Northumberland, Pennsylvania. Priestley showed that the gas he
had discovered was present in air and was closely connected with
burning. Let us trace the steps that led to the discovery of the true
nature of burning.
1) After his discovery of oxygen, Priestley, like a true scientist,
did not keep his discovery to himself. While he was in Paris later
on in 1774, he visited Lavoisier (la-vwa-zya') , the most eminent
chemist in France, and told him about his discovery. Priestley's in-
formation was a welcome addition to the many facts which Lavoi-
sier had already collected regarding the nature of burning.
2) Lavoisier examined carefully all the facts that he knew about
burning. He pondered over them for months, trying to formulate
an accurate theory that would explain burning and be in keeping
with all the observed facts.
3) Lavoisier was genius enough to use Priestley's work as the
basis of a theory which would explain the age-old puzzle of burning.
He suspected, and later advanced the theory, that when a substance
burned, it increased in weight because it united with something that
was present in air. Later this something was shown to be oxygen.
4) Other scientists had noticed this increase in weight when sub-
stances were burned in air. However, Lavoisier was the first to
formulate a theory based on this fact. He also undertook a series
of careful experiments to see whether their results would prove his
theory to be correct.
it
Lavoisier's classic 12-day experiment which explained burning.
"I introduced four ounces of pure mercury into a [sealed] glass ves-
sel," he wrote. "I lighted a fire in the furnace which I kept up con-
tinually for twelve days. On the second day, small red particles
already had begun to appear on the surface of the mercury." When
most of the mercury had been converted into a red powder, he re-
moved the glass vessel and its contents (which he had weighed before
the experiment) and weighed them again. There was no increase
in weight.
Since the glass vessel was sealed, nothing had entered or escaped
from it during the heating. Yet when he broke the seal he noticed
that air rushed into the vessel. To him this inrush of air indicated
that part of the air in the vessel had been used up during the heat-
ing, and had left space for more air to enter. After air had entered
the vessel, he weighed it once more and determined the increase in
weight. He concluded that this increase in weight equaled the weight
of something in the air in the vessel that must have combined with
the mercury, forming the red powder.
Lavoisier's inquiring spirit was not satisfied. He was a scientist in
the most modern sense. He refused to jump to a hasty conclusion on
the basis of a single experiment. He withheld drawing a conclusion
until he had performed many more experiments. As a further pre-
caution, he reversed his original experiment. He took the red pow-
der of mercury and heated it to a higher temperature. He found
that all of the red powder was changed back into mercury and that
a gas was given off, which he found by a series of tests to be identi-
cal with the oxygen gas that Priestley had discovered. Hence, he con-
cluded that it was the oxygen in the air that was responsible for
burning. Of all the substances he tried, he found none that could
burn without oxygen.
5) Burning, said Lavoisier, is the chemical union of a burnable
substance with oxygen. Simple enough. No mysterious phlogiston,
and the testimony of the most sensitive balance in Europe to sup-
port his reasoning. Thus, Lavoisier discovered the true explanation
of burning.
6) Lavoisier repeated the original experiment using other sub-
stances, including tin and sulfur. He found that the results of these
experiments were fully in accord with his theory. In this way his
theory was given further support.
14 NEW WORLD OF CHEMISTRY
"Nothing happens without a cause/9 said Leucippus 2500 years
ago. Ever since man first appeared on earth, he has been working
constantly to find out the why of many natural occurrences.
What are some of these occurrences that man has tried to explain?
Although these are only a few, we might include: Why does a stone
fall? What is fire? Why do some substances burn while other sub-
stances do not? What makes thunder? Why can birds fly?
In asking ourselves these questions and in answering them, we use
several words that probably were not used by primitive man. These
words, which appear later on in this paragraph, are words whose
meanings have come into rather common use comparatively recently,
as scientists measure time, perhaps within the last 25,000 or 50,000
years. Let us examine the first question. Probably we would ask:
What causes a stone to fall? By this we mean: What force causes a
stone to fall, for we know that a stone will not fall unless some lorce
acts on it, producing, we might say, an effect. We would answer the
question: A stone falls because the earth pulls the stone toward its
center (the force of gravity). As Leucippus (lu-slp'iis) implied,
modern scientists and most modern people believe that every cause
has an effect, and every effect has a cause. Such a relationship is
known as a cause-and-effect relationship.
Establishing a cause-and-effect relationship is not as simple as it
might seem. Early man did not have the many tools and instruments
which today we use so casually in finding cause-and-effect relation-
ships. Consequently, in attempting to explain the causes of a certain
effect, early man relied on what we would now consider magic, mys-
ticism, and superstition, but later on man learned to establish these
relationships by other methods. Then, too, it is sometimes difficult
to establish a cause, because often several causes taken together pro-
duce a single simple effect. Today many cause-and-effect relation-
ships are clearly understood; but on the other hand, the causes of
certain effects are not yet known, or even when known, are not thor-
oughly understood.
The method of deduction compared with induction. In explain-
ing a natural occurrence, Aristotle (ar'Is-tot"l) , a well-known teacher
and philosopher of ancient Greece, often made a bold and sweeping
general statement. From this general statement he drew inferences
and conclusions, which he thought applied in other similar cases.
Aristotle's method is commonly known as the method of deduction.
Francis Bacon, who lived almost 20 centuries later, was the first
man to make popular another method of reasoning. By a process con-
sisting of observation, collection of facts concerning the problem,
MATTER AND ITS CHANGES 15
formulation of a theory taking into account and explaining the
observed facts, and verification of the theory by actual experiment,
he formulated broad principles, sometimes called laws. Bacon's
method is known as the method of induction, and today it is used
widely. It is the pattern on which scientific method is based. To a
great extent, it is responsible for the success of scientific method.
While the method of induction is used very widely, deduction has
played its part too in the development of science. The establishment
of cause-and-effect relationships, usually by the method of induction,
is perhaps the greatest function of science.
What is scientific method? The method that Lavoisier used in
reaching the first correct explanation of burning is an example of
a pattern of action and thought used by scientists in their work. This
pattern is known as scientific method.
Scientific method may vary according to the nature of the problem
to be solved and the tools available for solving it. In general, how-
ever, the steps of the scientific method are represented by the six
steps you have just traced. In brief, they may be stated as follows:
The collection of all available facts related to a problem
The open • minded checking and examination of these facts
The formulation of a working theory based upon these facts
The testing of this working theory by experiments
The formulation of a law, or principle, from the tested theory
The use of this law in concrete and specific situations
tr ' '.
As you see, a scientific law, or principle, is a descriptive and ex-
planatory statement, or generalization, that expresses what men have
found to be accurate with respect to natural occurrences.
Triumph of scientific method. Lavoisier's explanation of burning
was not at once accepted. Indeed, for some time it met with bitter
opposition. Those who believed in the phlogiston theory attempted
to adjust their theory to fit the newly discovered facts, but this could
not be done. Even so, Lavoisier realized how hard it would be to
16 NEW WORLD OF CHEMISTRY
convince everyone of the truth of his own theory. He wrote: "I do
not expect that my ideas will be accepted at once; the human mind
inclines to one way of thinking, and those who have looked at na-
ture from a certain point of view during a part of their lives adopt
new ideas only with difficulty; it is for time, therefore, to confirm
or reject the opinions that I have advanced. Meanwhile, I see that
young men who are beginning to study the science (chemistry)
without prejudice or preconceived notions no longer believe in
phlogiston/'
New ideas, discoveries, and new theories as a rule must overcome
tradition and prejudice, but this should not discourage those who
introduce them. Tradition and prejudice are found not only in the
field of the natural sciences — physics, chemistry, astronomy, biol-
ogy, and others — but also in the social sciences — government, eco-
nomics, sociology, history, and others — in which they are likely to
be even more pronounced.
Lavoisier's theory of burning finally triumphed, however. The
accumulated evidence in its favor finally became so overpowering
that scientists could believe nothing else. The experimental method
of science had won over the strictly logical and theoretical method of
the ancients. Twenty years later Lavoisier, who was an aristocrat, was
beheaded during the frenzy of the French Revolution. "Until it is
realized that the gravest crime of the French Revolution was not the
execution of the king, but of Lavoisier, there is no right measure of
values, for Lavoisier was one of the greatest three or four men
France had produced." This statement, made by an eminent French-
man, expresses the judgment of thinking men the world over.
There are no authorities or dictators in science. There are only
those persons who know what men up to now have discovered; that
is, there are experts. In all of science there is no one who can say
"This theory is true." But there are many men who can say "On the
basis of what we now know, this theory seems to be sound." Scien-
tific method is a democratic method. In most cases, its theories are
the outcome of a pooling of facts discovered by many workers, a
cooperative effort.
Careful weighing and the law of the conservation of matter. To
Lavoisier a balance was absolutely necessary. By carefully weighing
all the substances entering into his experiments both before and
after each experiment, he found that there was no loss of weight
during the burning. The mercury plus the oxygen from the air
weighed exactly the same as the red powder of mercury (mercuric
oxide) . "One may take it for granted," he wrote, "that in every
MATTER AND ITS CHANGES
17
change there is an equal quantity of matter before and alter the op-
eration." In chemical changes we can change the form, the state, or
the composition of matter, but we cannot destroy matter itself. In
physical changes we can change the ionn or the state of matter, but
we cannot destroy the matter itself.
"But/* you might answer, "a candle burns until i( is all gone. It
becomes smaller and smaller and certainly weighs less and less."
And you are right. But if we take the trouble to collect and weigh
all the gases formed during the burning of the candle, we lind that
they weigh more than the original candle. This increase in weight
is due to the oxygen with which the burning candle has combined.
"What of an oak tree," you might say. "It grows from a little
acorn. Isn't matter created here?" It might seem so, but the fact that
many things increase gradually in size and weight does not mean
that matter has been or is being created. The oak tree does not come
from seed alone. The cells from which the tree is made are built up
chiefly from food materials taken out of the air, water, and soil.
Parts of the air, water, and soil have been chemically changed and
combined, forming living substances.
Matter can be neither created nor destroyed. This fundamental
law both of chemistry and of all science, is called the law of the con-
servation of matter.
1. Th« a* o
III* At
th« ttt«
th* the
tilti on ttt»
18
NEW WORLD OF CHEMISTRY
MATTER AND ITS CHANGES
19
Energy, too, can be neither created nor destroyed. So far we have
considered mainly the changes in matter that take place in burning.
But there are other changes that are equally important. When a
substance burns, heat is liberated. This heat may be used to convert
water in a boiler into steam. The pressure of this steam may then be
used to turn a wheel, thus producing rotary motion. By connecting
this wheel to a dynamo, electricity may be generated. This electricity
may be changed into heat or light or magnetism, depending on
whether it is sent through a toaster, a light bulb, or the electromag-
net of, let us say, a buzzer or telegraph sounder. Or this electricity
may be used to charge a storage battery. In charging the battery, the
electricity produces a chemical change, which, on being reversed,
yields electricity again.
Evidently all these — heat, electricity, the power to produce mo
tion, light, magnetism, and the power to produce a chemical change
— can be transformed one into the other. All are capable of doing
work, and all are forms of energy. In all energy changes, just as in
all changes in matter, there is no loss, only transformation. Energy
can be neither created nor destroyed. This is a fundamental law both
of chemistry and of all science. It is called the law of the conserva-
tion of energy.
Two laws or one? Researches on the structure of matter and the
nature of energy resulted in the atomic bomb. These researches,
one of
be of energy, follow
in the
discussed in Chapter 12, lead definitely to the conclusion that mat-
ter and energy are but different forms of the same thing, and that
matter can be converted into energy and energy into matter.
As a result, the law of the conservation of matter and the law of
the conservation of energy are no longer considered separate and
distinct laws. Instead, they may be considered as different phases of
a single law. Such a law would state that matter and energy can he
neither created nor destroyed, but that each can be transformed
into the other. In the transformation of matter into energy, matter
disappears and becomes energy. In the transformation of energy into
matter, energy disappears and becomes matter.
Can any form of energy produce a chemical change? When pa-
per is heated to its kindling temperature (see page 28) , it burns.
Heat, one form of energy, produces a chemical change. When light,
another form of energy, strikes a photographic film, it causes a chemi-
cal change in the substances that coat the film. Thus, light also can
cause a chemical change. When an electric current is passed through
water, it splits the water into two gases, neither of which resembles
water vapor. Thus electricity, too, is very effective in producing a
chemical change. (The energy that was used to split the water ap-
pears again as heat energy when the two gases are recombined, form-
ing water.) From these experiments and from others, we know that
many forms of energy bring about chemical changes.
" '*' "*•'- """~
-f*
into-
**-
20
NEW WORLD OF CHEMISTRY
YOU WILL ENJOY READING
The literature of chemistry is filled with romance. In it you
Will find stories of human struggle and achievement, stories
whose truth makes them all the more worthy to be read and
remembered. Even the most interesting and significant parts
of this literature cannot all be listed in this book. Neverthe-
less, they are available, and it will be well worth your while
to read as many of them as you can. The following is a selected
list of books that deal with some of the topics in this chapter.
Becker, Carl; Painter, Sidney; and Han, Yu-Shan. The Past
that Lives Today, pp. 22-86. Silver Burdett Co., New York,
1952. A fascinating story of early man, ancient civilizations
and the science of the ancients is told here.
Fabre, Jean H. The Wonder Book of Chemistry, pp. 6-69.
Albert & Charles Boni, New York, 1922. A delightful account
of elements, mixtures, and compounds.
Jaffe, Bernard. Crucibles: The.Story of Chemistry, pp. 34-50.
Simon and Schuster, New York, 1948. A simple account of the
phlogiston theory.
Somerville, John. The Way of Science: Its Growth and
Method, pp. 93-113. Henry Schmnan, New York, 1953. A very
simple illustration ot the steps ot scientific method.
USEFUL IDEAS DEVELOPED
1. The three states of matter are solid, liquid, and gaseous.
2. In a physical change, the original substance does not
change into one or more other substances.
3. In a chemical change, the original substances disappear
and new substances are formed.
4. An element is a substance that cannot be broken down,
or decomposed, into a simpler substance by the ordinary types
of chemical change.
5. A chemical compound is a substance composed of two
or more elements so combined that the elements can no longer
be identified by their original individual properties.
6. A mixture is a substance composed of two or more ele-
ments or compounds that still retain their individual prop-
erties after they have been thoroughly mixed.
7. Burning is the chemical change in which a burnable
substance unites with oxygen.
8. The law of the conservation of matter states that mat-
ter can be neither created nor destroyed. But matter may be
changed from one form to another.
MATTER AND ITS CHANGES 21
9. The law of the conservation of energy states that energy
can be neither created nor destroyed. But energy may be
changed from one form to another.
10. Recent researches prove that matter can be transformed
into energy, and that energy can be transformed into matter.
11. Establishing accurate cause-and-effect relationships is
perhaps the greatest function of science.
12. The steps in the scientific method include (1) the col-
lection of all available facts related to a problem, (2) the
open-minded examination of these facts, (3) the formulation
of a working theory based upon these facts, (4) the testing of
this working theory by experiments, (5) the formulation of a
law, or principle, from the tested theory, and (6) the use of
the law in specific situations.
13. Blind acceptance of so-called "authorities," prejudices,
and personal likes and dislikes have no place in the general
pattern of action and thought of the true scientist.
USING WHAT YOU HAVE LEARNED
Group A
1. (a) Which of the following are physical changes and
which are chemical changes? (b) Give reasons for your
answers. Souring of milk, molding of clay, digestion of food,
drying of clothes, dissolving sugar in water, decay of fruit,
freezing of mercury, photosynthesis (manufacture of starch in
the leaves of plants from carbon dioxide gas and water) ,
erosion (the breaking up of rocks and soil by the action of air
and water) .
2. (a) What were the "four elements" of the ancients?
(b) In what respects did their observations seem to support
this theory? (c) What was one weakness of the way in which
this theory was formulated?
3. What are your reasons for thinking that water is a
compound?
4. Classify the following as elements, compounds, or mix-
tures: table salt, mercury, aluminum, paper, carbon dioxide,
gold, silver, iron rust, sugar, sulfur, milk, brass, silver coin.
5. Name a compound that contains hydrogen.
4. Explain one method of telling a mixture from a com-
pound. «•
22 NEW WORLD OF CHEMISTRY
7. Give briefly the main ideas of the "phlogiston theory."
8. What was Priestley's part in discovering the true nature
of burning?
9. (a) Complete the following statement: When red mer-
curic oxide is heated, it is changed into and (b) Is
the change a physical or a chemical change? (c) Why?
I . T .
' i
10. Describe Lavoisier's 12-day experiment on burning.
11. Is Lavoisier's explanation the modern explanation of
burning?
12. What part did the balance play in the development of
chemistiy?
13. State the six general steps in the scientific method.
14. Distinguish between the method of science and the
method ol the ancient teachers and philosophers.
I ...
15. (a) What is the law of the conservation of matter?
(b) Is the disappearance of camphor balls in clothes an excep-
tion to this law? (c) Explain.
16. (a) What is energy? (b) Name three different forms of
encigy.
17. Give an illustration to show that each of the forms of
eneigy named in your answer to question 16 may produce a
chemical change.
18. Assume that coal is the source of energy that lights your
home (electric light). Make a list of the transformations of
energy that occur, beginning with the burning of coal and end-
ing with the lighted bulb.
Group B
19. Not many of the chemical elements wrere discovered by
Americans, (a) Which elements are these? (b) Can you sug-
gest a reason why Americans have discovered so few?
20. Aluminum is the most abundant metal in the earth.
Tell why only in recent years aluminum has come into com-
mon use.
21. (a) Can one use scientific methods in fields other than
science? (b) Explain your answer.
22. (a) State four evidences of chemical action, and (b) give
one example of each.
MATTER AND ITS CHANGES 23
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Discuss with your teacher of history the problem of
prejudice and tradition as obstacles to progress. Write a report
on this topic using an illustration from American history or
from the history of science.
2. Find an article on some subject such as mental telepathy,
communications from the spirit world, extrasensory percep-
tion, astrology. Read the article carefully and write your own
reaction as to whether you, using the methods of science,
would accept the author's conclusions as scientific.
3. Obtain a new photoflash bulb and weigh it on a sensitive
balance. Weigh it again after it has been ignited. Do your
findings uphold the law of the conservation ol matter?
4. Show how you would use the method of science in solving
some particular everyday problem which you have had to
solve or which you will soon have to solve.
5. With the aid of a medicine dropper, allow a drop of ink
to fall into a tall glass of cold water. Observe what happens
after approximately 5 minutes, 8 hours, 2-1 hours, and 2 days.
Report to class with reference to physical and chemical
changes, and any other conclusion you have drawn.
6. Prepare a brief report on an expert in some field of sci-
ence. Find out how he became an expert. How does an expert
differ from an authority?
7. On the first page of this chapter, there is a statement
made by Lavoisier. Read it carefully. Report to your class on
the importance of this statement. What are its implications
for everyday living?
EARTH'S MOST ABUNDANT ELEMENT
. . . / procured a mouse and put it
into a glass vessel containing the air
(oxygen) from the red powder of
mercury. Had it been common air, a
full-grown mouse, as this was, would
have lived in it about a quarter of an
hour. In this air, however, my mouse
lived a full half hour.
Joseph Priestley, 1775
A Sunday experiment by an English minister. On Sunday, the
first of August, 1774, Priestley was working in his laboratory. He
placed a red powder (mercuric oxide) in a bell jar so arranged that
any gas which might be formed would pass out of the bell jar
through a tube and be collected in a bottle.
Instead of heating the powder over a flame, he used a large burn-
ing lens to concentrate the rays of the sun on the powder. "I pres-
ently found/' lie reported, "that air was expelled from it readily."
But this result was not startling, because others before him had ob-
tained gases by heating solids.
A candle was burning in the laboratory. Wondering what effect
this gas would have on a flame, he placed the candle in a bottle of
it. Priestley reported in somewhat flowery words: "A candle burned
in this air with an amazing strength of flame; and a bit of red-hot
wood crackled and burned with a prodigious rapidity exhibiting an
appearance something like that of iron glowing with a white heat
and throwing out sparks in all directions.'*
Only natural curiosity or perhaps chance led Priestley to experi-
ment with the gas (or as he called it, air) . As Priestley, himself, said
later, he had no idea what the outcome might be.
OXYGEN 25
Priestley was unable to explain what had happened. He was such
a firm believer in the phlogiston theory that he did not associate this
new gas with burning. As we learned in Chapter 1, it was Lavoisier
who showed that Priestley's air (later called oxygen) is the element
necessary for burning, thus solving a mystery that had baffled scien-
tists for centuries.
Half the earth is oxygen! Priestley's discovery of oxygen was a
turning point in the development of chemistry. It is one of the
strangest facts in history that this element, which surrounds us every-
where and without which life is impossible, was not obtained pure
until about 180 years ago. This fact is even more surprising when
we realize that this one element, oxygen, is present on earth in quan-
tities equal to the weight of all the other chemical elements put
together. Sand and half of the different kinds of rocks on the earth
are compounds of oxygen. Water contains almost 99 percent oxygen
by weight, and air contains about 21 percent oxygen by volume.
How oxygen is prepared in the laboratory. Because mercuric
oxide, from which Priestley prepared oxygen, costs about one dol-
lar a pound, it is too expensive to use in the laboratory. Instead, we
obtain oxygen from a white crystalline solid called potassium chlo-
rate. This compound is composed of three elements — potassium,
chlorine, and oxygen. By applying heat, the oxygen can be torn
away and liberated as free oxygen.
How to express the change by which oxygen is prepared. As po-
tassium chlorate is heated, it yields potassium chloride plus oxygen.
This is a chemical reaction which produces the chemical change rep-
resented by the equation:
Potassium (K) \ -\
^, , . /X«; potassium
Chlorine (Cl) \ -*> rul . > -f oxygen
,lv I chlorine J 75
Oxygen (O) J J
2KC1O3 -» 2KC1 + 302
potassium chlorate — > potassium chloride -f oxygen
The chemical shorthand used to express the change will be fully
explained later. The forms of energy that take part in the change are,
as a rule, omitted from the equation. A chemical change is the re-
sult of a chemical reaction.
How the speed of this reaction can be increased. The method of
preparing oxygen just described has one serious drawback. Unless
a very high temperature is reached, oxygen is liberated very slowly.
Someone discovered that if a small amount of powdered manganese
dioxide, a black solid, is added to the potassium chlorate before
26 NEW WORLD OF CHEMISTRY
heating, oxygen is liberated more quickly and at a lower tempera-
ture.
At the end of the chemical reaction, that is, when oxygen is no
longer given off, the same amount of manganese dioxide with which
the experiment started remains. The weight of manganese dioxide
has not been changed in any way. None of its oxygen has been lib-
erated. Since manganese dioxide remains unchanged at the end of
the reaction, it is not included in the equation.
A catalyst changes the speed of a chemical reaction. Chemists
have discovered that many chemical reactions can be speeded up or
slowed down by placing a small quantity of certain substances in
contact with the reacting materials. A substance that changes the
speed of a chemical reaction is called a catalyst, or a catalytic agent.
The catalyst itself may undergo some temporary change, but at the
end of the reaction is present in the same state and quantity as at
the beginning. Manganese dioxide, as used in the laboratory prepa-
ration of oxygen, is a catalyst. However, manganese dioxide does
not always act as a catalytic agent. It (see page 137) actually enters
into certain reactions in which its composition is changed perma-
nently.
We know little about the reasons that catalysts act as they do.
We do know, however, that many vitally important chemical re-
actions take place too slowly, and hence uneconomically, except in
the presence of certain catalysts. Research on the nature of catalysis
is now being carried on in laboratories throughout the world, and we
will know much more about it before very long.
The presence of a catalyst in a reaction is sometimes indicated by
writing the catalyst over the arrow in the equation. In the reaction
just discussed, the presence of the catalyst, manganese dioxide, might
be indicated thus:
MnO2
2KC1O3 > 2KC1 + 3O2
How oxygen gas is collected. In preparing oxygen, a mixture of
potassium chlorate and manganese dioxide is put in a test tube and
heated over a bunsen burner (see illustration page 27) . Connected
to the test tube is a delivery tube which reaches into a bottle that
has been filled with water and placed mouth downward in a pan of
water. As the potassium chlorate is heated, it is broken down, or
decomposed, forming potassium chloride and oxygen.
The potassium chloride is a solid and remains in the test tube.
The oxygen is a gas and passes through the delivery tube into the
2. of if
ii the
Why?
water-filled collecting bottle from which it displaces water. This
method of collecting gases is called the displacement of water method,
and was used by Priestley. Priestley also collected gases by the dis-
placement of mercury when such gases dissolved in water but not
in mercury. If a gas is collected by the displacement of a liquid,
chemists say that the gas is collected over the liquid. Thus, Priestley
collected oxygen over water.
Physical properties of oxygen. In discussing the physical properties
of a gas, we usually consider five characteristics: (1) color, (2) odor,
(3) weight compared with the weight of an equal volume of air,
(4) the ease with which it may be changed into a liquid, and (5) its
absorption by water. This absorption by water we call its solubility
in water. Just as solids, such as sugar and salt, disappear as solids
when stirred in water and are distributed uniformly throughout
the liquid, so gases, such as air and oxygen, when passed through
water, may likewise be absorbed by the water. In the case of some
gases, taste is considered also.
Oxygen is colorless and has no odor. It is slightly heavier than
air. It is slightly soluble in water. Under normal conditions, that
is, at a temperature of 18 degrees centigrade and a pressure of one
atmosphere, or 760 millimeters of mercury,* about four quarts of
* Scientists often use two terms to describe temperature and pressure, normal
conditions and standard conditions. Standard conditions mean a temperature
of 0° C. and a pressure of 760 mm. of mercury, that is, one atmosphere (atm.) .
Normal conditions mean, in effect, roorn temperature and the pressure of 1 at-
mosphere. Throughout this book a pressure of 760 mm. is assumed when staling
temperatures at which gases liquefy or liquids solidify. A millimeter (mm.) is
a small unit of length in the Metric System. Numerically it is equal to 0.001 of
a meter, a larger unit of length equal to 39.37 inches. For an explanation of
metric units and temperature scales, see pages 654 and 644 respectively. Mcttic.
units of measurement are used by scientists in all countries.
28 NEW WORLD OF CHEMISTRY
oxygen gas will dissolve in 100 quarts of water. It is hard to change
oxygen to a liquid. At a temperature of about 183 degrees below-
zero centigrade (— 183°C.) and under a pressure of 760 millimeters,
oxygen is converted into a pale blue liquid, which can be attracted
by a magnet. At — 219°C. it changes into a bluish-white solid.
Chemically, oxygen is an active element. Because oxygen combines
with almost all other elements, forming compounds called oxides, it
is considered to be very active chemically. For example, when iron
is exposed to oxygen, rust, which is an oxide of iron, is formed.
Iron + oxygen — > iron oxide
4Fc + 3O2 -> 2Fc2O3
If iron is first heated until it glows and then placed in a bottle of
oxygen, the chemical reaction is so vigorous that the iron burns
brilliantly, throwing off sparks of glowing iron oxide. This sur-
prising spectacle actually is iron burning in oxygen, just as a piece
of paper burns in air!
Differences between slow and rapid oxidation. When a sub-
stance— element or compound — combines with oxygen, new sub-
stances are formed. This chemical union of a substance with oxygen
is called oxidation. Rapid oxidation, such as the burning of coal, is
accompanied by noticeable heat and light. During slow oxidation,
such as the rusting of iron or the decaying of wood, no light is given
off nor can we easily detect the heat because it is given off so slowly.
Delicate measurements, however, have proved beyond doubt that
the amount of heat energy liberated is the same whether the oxida-
tion of a substance takes place slowly or rapidly.
The soft cold light of the firefly, and the glow of some fungi and
bacteria are caused by the oxidation of a complex chemical com-
pound, luciferin, which they produce.
Why some substances catch fire more easily than others. Some
substances catch fire at low temperatures, but others require ex-
tremely high temperatures in order to burn. Every substance must
be raised to a certain definite temperature before it can combine
with oxygen at such a rate that the heat produced is sufficient to keep
the substance burning without the addition of more external heat.
In starting a coal fire, we often begin by burning paper, which
sets fire to kindling wood, which sets fire to the coal. The heat given
off by the burning paper causes the wood to catch fire; the heat given
off by the burning wood in turn causes the coal to catch fire. The
lowest temperature at which a substance catches fire and continues
to burn is called the kindling temperature of that substance. We
Hitreau of Mintt, U.S. Depf. of Intu
A coal-dust explosion issues from the U.S. Bureau of Mines*
Experimental Coal Mine in Pennsylvania. Such test ex-
plosions are part of the Bureau's safety research program.
say that a substance which is easy to set on fire has a low kindling
temperature, and a substance which is difficult to set on fire has a
high kindling temperature.
A substance may have different kindling temperatures, depending
upon the size of the particles into which it is divided, that is, upon
its state of subdivision. A solid piece of iron has a high kindling
temperature, but powdered iron, because of the large surface that
is exposed to the oxygen of the air, can be made to burn readily
in air. Many dust explosions in flour mills, starch factories, grain
elevators, and coal mines are caused by the very rapid oxidation of
explosive mixtures of air and finely divided materials. A spark
resulting from static electricity or friction often sets off the explosion.
Smut dust and air form an explosive mixture which may be
ignited by static electricity during threshing operations. Costly fires
of this kind have been so widespread that the United States Depart-
ment of Agriculture has issued a bulletin explaining how to prevent
them.
30 NEW WORLD OF CHEMISTRY
The explosion of a mixture of coal dust and air has been used
in one type of internal-combustion engine.
Substances such as asbestos, brick, concrete, and marble never
catch fire because they are already completely oxidized.
Spontaneous combustion. Several years ago there were widespread
floods in the Ohio Valley. The lower parts of thousands of haystacks
in the Valley were soaked with water. As the flood waters receded,
farmers were pu//led when some of the haystacks began to catch fire.
The explanation of this phenomenon is simple. During respiration
of living plant cells, food materials slowly oxidize and heat is given
off. This oxidation is speeded up by the presence of a small amount
of moisture. The hay itself also slowly oxidi/es, liberating heat. As
the temperature rises, the rate of oxidation also increases. The heat
slowly accumulates, and when the kindling temperature of the hay is
finally reached, it bursts into flame. Materials catching fire in this
way are said to undergo spontaneous combustion. (Combustion
refers to any chemical reaction which produces heat and light.
Burning is only one kind of combustion.)
Fires have also been caused by painters' rags saturated with linseed
oil. As the linseed oil slowly oxidizes, heat is given off. Unless there
is a free circulation of air to carry away this heat, the oily rags may
become hot enough to catch fire. Thus, you see why oily rags should
not be kept in poorly ventilated closets. Finely divided coal in the
closed hold of a ship, or in a poorly ventilated boiler room must be
sprinkled with plenty of water from time to time to prevent the
accumulation of heat from slow oxidation.
Phosphorus, an element that burns spontaneously. Ancient alche-
mists spent most of their time looking for the philosopher s stone,
which they believed would change lead into gold. In 1669, while
searching for the philosopher's stone, Hennig Brand, an alchemist
of Hamburg, obtained a new and strange chemical substance from
urine. It had so low a kindling temperature that on exposure to
air it caught fire immediately and burned, forming a white powder
(phosphorus oxide) . Brand made a famous tour of Europe, exhibit-
ing this unusual substance. Today we know it as white phosphorus,
a soft, waxy element, now obtained by a chemical process from bone
deposits. At one time it was used in the manufacture of matches
(see page 296) . This element is sometimes referred to as yellow
phosphorus.
To keep white phosphorus from catching fire, it must be stored
under water. Do not touch it with your bare fingers, for white
phosphorus will cause severe burns that heal very slowly.
OXYGEN
31
Upon exposure to air, white
phosphorus ignites sponta-
neously in a violent reaction.
Note the protective clothing
worn by the demonstrator.
Monsanto Chemical
You could not live without oxygen! One of the remarkable ex-
periments which i'riestley performed showed that a mouse, placed
in a bell jar, lived twice as long in pure oxygen as in "common air."
This, Priestley was unable to explain; but today we know that the
chief chemical change that goes on in the body of any living animal
is slow oxidation.
Man obtains his supply of oxygen by breathing. When he inhales,
oxygen is taken into his lungs. This oxygen passes through the walls
of the lungs and is absorbed by the red cells of the blood, which
carry it to all parts of the body. In every living cell or body dim,
oxidation takes place, liberating heat and other forms of energy. This
slow, steady oxidation is like a tiny flame which keeps life going.
Without oxygen, the flame is snuffed out and life is extinguished.
With the exception of a few very low forms of life, all living
things take oxygen from the air. This oxygen is in the /rvv, or uiic.oin-
bined, state. That is, the oxygen is not chemically united with any
of the other substances in air. Fish obtain their supply ol oxygen,
from the air that is dissolved in water.
stem
Adapted from drawing by Linde Air Products Company
Fig. 3. Oxyacetylene torch. What is the
function of the expansion chamber?
Industry uses great quantities of oxygen. Commercial production
of pure oxygen in the United States is more than 25 billion cubic
feet a year. Of the oxygen produced, it is estimated that more than
95 percent is used in cutting and processing steel and in welding
metals, such as aluminum and steel, by means of oxyacetylene and
oxy hydrogen torches.
It was Priestley who first thought of using oxygen to produce high
temperatures. He found that blowing pure oxygen on a piece of
glowing wood would cause it to burn furiously. A few years later
an American scientist, Robert Hare, of Philadelphia, put this dis-
covery of his friend Priestley to practical use by inventing the oxyhy-
drogen torch, or blowpipe.
The oxyhydrogen torch consists of two tubes, one inside the other.
Hydrogen gas passes through the outer tube and is ignited at the tip
of the torch. Pure oxygen passes through the inner tube and the
mixture of the two gases burns at the tip of the torch with an ex-
tremely hot flame, about 2400°C., a temperature much higher than
the melting point of iron.
The oxyhydrogen torch was never widely used. Instead, the oxy-
acetylene torch is used. With the oxyacetylene torch, a flame tem-
perature of over 3300°C. may be easily produced. The oxyacetylene
torch is similar in principle to the oxyhydrogen torch, but acetylene
Fig. 4. Blast lamp. Similar in principle
to the oxyhydrogen torch except
that compressed air and any fuel
gas are used instead of oxygen and
hydrogen. The lamp is used by glass
blowers and jewelers.
flame
air
gas
OXYGEN
33
gas is used instead of hydrogen (see Fig. 3) . The oxygen used in
the torch is stored under high pressure in strong steel cylinders.
The acetylene, however, is not under high pressure, but is dissolved
in a liquid called acetone. In almost every automobile service garage,
oxyacetylene torches, with their accompanying cylinders ol oxygen
and acetylene, may be seen ready for use.
The oxyacetylene torch is an important industrial tool. It is used
to weld, cut, and clean metal. It is also used in heat treating sur-
faces of metal machine parts to make them more wear resistant.
Oxygen rusts and derusts steel. In the presence of air and minute
quantities of water vapor, steel rusts. In rusting, oxygen from the
air unites slowly with the metal, forming a brown, scaly oxide of
iron. At the high temperatures used in the making of steel, rust lorms
very rapidly and is a serious problem. As red-hot steel is carried to
the rolling mills, it becomes covered with seams of iron oxide.
These surface imperfections are removed by a process called
torch-deseaming, or scarfing. In the scarfing operation, the llame of
the oxyacetylene torch is directed onto the hot steel. The surface
of the steel is quickly oxidized to a depth of about one quarter of an
inch. The iron oxide falls off readily, leaving an unblemished sur-
face. Scarfing may be done by hand or by special machines.
(left) Red-hot steel slab passes through scarfing machine
in which oxyacetylene torches remove surface defects,
(right) Welding metal plates with an oxyacetylene torch.
Linde Air Products Company
34
NEW WORLD OF CHEMISTRY
Because of the high tem-
perature of the torch, an
oxyacetylene weld is
smooth and strong.
/.nidc Air Product* Company
Oxygen saves lives. Priestley also discovered another use for oxy-
gen. Alter inhaling oxygen, he wrote: "My breath felt peculiarly
light and easy. It (oxygen) may be peculiarly salutary to the lungs in
certain cases where the common air is not sufficient."
Today, pure oxygen is administered to persons with pneumonia
and in other cases where the respiratory system cannot function at
its normal rate. Usually about two gallons of oxygen are administered
per minute. Since air contains only about 20 percent pure oxygen,
a patient who is weak, or whose lungs are congested or partially
destroyed, can satisfy the oxygen requirements of his body by breath-
ing a much smaller volume of air that is rich in oxygen than he
would normally require of ordinary air. Oxygen is usually admin-
istered by means of an oxygen tent, a canopy which fits over the
patient's bed. Pure oxygen is introduced at such a rate that the air
inside the canopy always contains from 45 to GO percent oxygen.
Oxygen in determining basal metabolism. Pure oxygen is used by
physicians in determining the rate at which a person's food supply
is oxidized while the person is at rest. This is known as his rate of
basal metabolism. This rate is obtained by measuring the volume of
oxygen consumed by the person at rest during a short interval,
usually eight minutes.
From these data the number of liters* of oxygen consumed per
minute may be calculated easily. This number is then compared
with the basal metabolic rate for a normal person of the same age,
sex, height, and weight. Persons in normal health use oxygen at a
* A liter (I.) is a unit of capacity (volume) in the Metric System. It is
slightly larger than a U.S. liquid quart.
OXYGEN
35
standard rate. In certain diseases, the patient's rate is higher and
in others it is lower than known standards for persons in good
health. For example, a high basal metabolic rate always accompanies
an overactive thyroid (see pages 147, 581). A low basal metabolic
rate may be an indication of an underactive thyroid.
Oxygen flies high. Aviators and mountain climbers who ascend to
high altitudes where the atmosphere is very thin must carry supplies
of oxygen. Otherwise their senses become dulled and they are likely
to lose consciousness. The United States Air Force requires the use
of oxygen at altitudes above 10,000 feet in the daytime and from the
ground up at nighttime. Airliners flying at altitudes ranging from
15,000 to more than .SO, 000 feet contain equipment to keep the oxy-
gen concentration inside their cabins only slightly less than that at sea
level. The cabins are airtight and are pressurized by means of pumps,
so that the inside pressure does not become uncomfortably low.
Rescue parties entering mines and buildings in which dangerous
gases are present carry oxygen-breathing apparatus.
Air Photographic and Charting Service, U.S. Air Force
High in the strato-
sphere, this pilot is de-
pendent upon a con-
tinuous supply of pure
oxygen.
36 NEW WORLD OF CHEMISTRY
Some other uses of oxygen. Oxygen is used in the photoflash lamps
employed in photography. These lamps look like ordinary electric-
light bulbs, but they are filled with oxygen and aluminum foil. When
an electric current is passed through the filament, the aluminum foil
is raised to its kindling temperature. It ignites and oxidizes with
a blinding flash. This Hash lasts only about ^ of a second.
Pure oxygen has also been introduced in the manufacture of steel,
synthetic gasoline, and fuel from underground coal (see pages 335,
3(50, 417).
How is such a huge amount of oxygen prepared? It is not sur-
prising to find that the two main sources of oxygen are air, the most
abundant mixture containing oxygen, and water, the most abundant
compound ol oxygen. Nearly all commercial oxygen is obtained
from liquid air by first lowering the temperature of the air until it is
changed to a liquid. The preparation of oxygen from liquid air is
discussed on page 100. Less than one percent of the oxygen produced
commercially is obtained by decomposing water by means ol an
electric current. This process, called the electrolysis of water, is de-
scribed on page 62.
How can we test for oxygen? When Priestley prepared oxygen
from mercuric oxide, he tested the gas by placing burning and
glowing substances in it. In each case, the substance burned more
vigorously. This method is still used to identity oxygen. A glowing
splint or splinter ol wood is thrust into a bottle oi the gas. Such a
splint placed in a bottle of oxygen bursts into flame at once.
No other odorless gas will cause a glowing splint to burst into
flame in this way. Hence, we can distinguish oxygen from any other
odorless gas by this simple procedure. We call such a method of
identifying a substance a test for that substance. Chemists have
devised hundreds of tests, which they use in identifying many other
pure substances.
Ozone, the active. Ten years after Priestley's discovery of oxygen,
another gas which possessed a peculiar odor and which, unlike
oxygen, tarnished mercury under normal conditions, was reported.
But it was not until 1810 that Schoenbcin (shun'bfn) isolated this
gas and called it ozone, from the Greek word meaning to smell. Its
sharp odor is noticeable around electric machines in operation.
O/one is a pale blue gas, one and one-half times as heavy as
oxygen. It is even less soluble in water than oxygen but is more
active chemically. It is a strong oxidizing agent. That is, it is a
substance which readily supplies oxygen for chemical union with
another substance.
/. -'
current L'"11'
Fig. 5. A continuous-process ozone tube. Dry air enters
at lower left, passes through the brush discharge of high-
voltage current. Part of the oxygen of the air is converted
to ozone which leaves through pipe at right.
Ozone (written O8) is prepared by passing electric discharges
through cither dry air or oxygen (written O.,) . About eight percent
of the oxygen is converted into pure o/.one, although slightly larger
yields are obtained if the temperature is kept low, or if the process
is continuous.
3 volumes of oxygen
302
- 2 volumes of ozone
203
Ozone is unstable; it changes back to oxygen quickly, two volumes
of ozone changing into three volumes of oxygen. It cannot be stored
and must be produced at its point of use.
What is allotropy? As we have just learned, oxygen exists in two
forms: ordinary oxygen and o/one. The existence in the same
physical state (both oxygen and o/.one are gases) of two or more
forms of the same element is a phenomenon called allotropy
(0-lot'r6-pi) . The various allotropic forms of an element have differ-
ent physical and chemical properties.
The cause of allotropy is not yet completely understood. Hut we
know that it is caused in part by differences in the arrangement of
the atoms and in the amount of energy in the various allotropic
forms. This can be seen easily by referring to the way in which o/one
is produced from oxygen. When an electric current discharges
through oxygen, the electric energy changes oxygen into ozone,
which, as you would expect, possesses more energy than ordinary
oxygen. O/one, on changing back into oxygen, liberates this energy
in the form of heat. During the change from oxygen to ozone, or
from o/one to oxygen, no energy is destroyed, nor is any energy cre-
ated. The change from one allotropic form of oxygen to the other
is an excellent example of the law of the conservation of energy.
37
38 NEW WORLD OF CHEMISTRY
Allotropy is not confined, of course, to oxygen. For example, the
element phosphorus occurs in two allotropic forms: white phos-
phorus and red phosphorus. White phosphorus melts at 44°C., is
poisonous, soluble in carbon disulfide, and has a very low kindling
temperature. Red phosphorus is nonpoisonous, insoluble in carbon
disulfide, and has a higher kindling temperature than white phos-
phorus. Red phosphorus is also heavier and less active chemically
than white phosphorus. Under proper conditions, red phosphorus
may be changed into white phosphorus, and vice versa.
Ozone "burns up" germs. Because of its extreme chemical activity,
o/one is used to a limited extent in purifying water. It kills bacteria
and other microorganisms in water by oxidi/ing them, literally burn-
ing them up. Other organic: materials present are also oxidized.
About one gram* of ozone will purify a cubic meter of water.
Ozone destroys odors. Because o/one is an excellent oxidi/ing
agent, it is used also in purifying air in homes, refrigerators, tunnels,
and zoos. Small ultraviolet lamps change the oxygen in air to ozone
which clears away bad odors. Because of the increasing abundance of
low-cost electricity, it seems possible that o/one may be used more
widely in the removal ot unpleasant tastes and odors from water
than it now is.
Ozone helps screen the earth. Ozone is present in the layers of
the atmosphere, about 30 miles above the surface of the earth. Sci-
entists believe that this region which is rich in o/one acts as a screen
that protects life on earth from the harmful effects of too much ultra-
violet light from the sun.
What part does chance play in scientific discoveries? Some years
after his discovery of oxygen, Priestley commenting on this memo-
rable occasion said: "I can not at this distance of time recall what it
was that I had in view in making this experiment, but I had no
expectation of the real issue of it. If I had not happened to have had
a lighted candle before me, I should probably never have made the
trial, and the whole train of my future experiments relating to this
kind of air might have been prevented. More is owing to what we
call chance than to any proper design or preconceived theory."
Chance may have played some small part in leading Priestley to
make his experiments. It seems likely, though, that he failed to take
into account the consuming natural curiosity, always present in true
scientists, which literally forced him to make his experiments. No
* A gram (g.) is a small unit of weight or mass in the Metric System. One
thousand grams are equal to a kilogram, which weighs slightly more than 2.2
pounds. For a definition of mass, see page 157.
OXYGEN 39
doubt one of the tests he would eventually have made on any gas
was to see if it would burn. If the lighted candle had not been
present at that particular moment, he undoubtedly would have
lighted one later on. Probably what Priestley meant by the statement
was simply that he had no specific purpose in mind in making the
experiment. But having no specific purpose in mind and attributing
the results to chance are not the same. Nevertheless, such confusion
exists among scientists even today.
Other men were performing experiments similar to those made by
Priestley (among them Scheele) , and we might say that the time was
ripe for these discoveries. If Priestley had failed to make them, no
doubt some other experimenter soon would have made them.
Today carefully planned research has speeded up advances in
science. Chance now plays a very slight role in the development of
chemistry. Surely we cannot regard as "chance" the keenness of
mind to appreciate the significance ol, and to follow up by intelligent
experiment, a clue furnished by some unforeseen event.
Progress in science and changes in society are closely interrelated.
Most of us, although we live in a world ol science, have strange
notions about what scientists arc like. We picture the research man
as a lone, mysterious genius who locks himsell up in his laboratory
away from the world, and attempts to solve some abstract scientific
problem. We sec him emerging triumphant, after weeks or perhaps
years, with some great discovery. Most of us have an idea that the
efforts of the man of science arc influenced very little by the society
in which he lives. Nothing could be farther from the truth. The
scientist is influenced by society and society in turn is influenced
by the scientist.
The steam engine, for example, was developed out of the social
needs of the eighteenth century. With the Industrial Revolution,
which began in Birmingham, England, came a need for more iron to
make the machinery so much in demand. In making iron, wood char-
coal had been used, but England's forest reserves had been seriously
depleted by the use of timber in the many ships swallowed up in
naval wars and commercial ventures. As a result, coal had to be sub-
stituted for charcoal, and coal mines, abandoned because they were
Hooded, had to be reopened. This meant draining mines, and James
Watt improved Newcomen's steam engine for this purpose.
At about the same time, the problem of burning was reinvesti-
gated, and Priestley's investigations (which led to his discovery of
oxygen) were closely related to society's need for more information
concerning the extraction of metals.
40
NEW WORLD OF CHEMISTRY
YOU WILL ENJOY READING
Becker, Carl. Modern History, pp. 581-607. Silver Burdett
Co., New York, 1952. "How Science Gave Men Machines to
Work for Them, and How the Machines Changed the Con-
ditions under Which Men Had to Live and Labor." Very
readable.
Ficklen, Joseph B. "Dust Explosions." Journal of Chemical
Education, 'March, 1942, pp. 131-134. Published by the Amer-
ican Chemical Society, Easton, Pa. Editorial office: Metcalf
Chemical Laboratory, Brown University, Providence, R.I. An
interesting and well-illustrated article.
Friend, J. N. Man and the Chemical Elements. Scribners,
New York, 1953. The author rambles along into many inter-
esting sidepaths and offers sidelights such as the origin of
gibberish from the alchemist Geber.
Partington, J. R. A Short History of Chemistry, pp. 110-120.
The Macmillan Co., London, 1939. The fascinating story of
Priestley's discovery of oxygen.
USEFUL IDEAS DEVELOPED
1. A catalyst is a substance that changes the speed of a
chemical reaction without being itself permanently changed.
2. Oxidation is the chemical union of a substance with
oxygen. In slow oxidation, neither light nor noticeable heat is
liberated. In rapid oxidation, light and noticeable heat are
evolved. Burning is rapid oxidation.
3. The kindling temperature ot a substance is the lowest
temperature at which that substance catches fire and con-
tinues to burn.
4. Spontaneous combustion is a burning started by the heat
accumulated during slow oxidation. Combustion is any chem-
ical action that liberates heat and light.
5. Finely divided powders form explosive mixtures with air
because of the extremely large surfaces exposed to oxygen.
6. Elements may occur in two or more varieties, or allotropic
forms, in the same physical state. These allotropic forms differ
in their properties because of differences in the arrangement
of their atoms and differences in the amount of energy pos-
sessed by particles of each allotrope.
7. Today carefully planned research has speeded up ad-
vances in science, and chance discoveries play a smaller part
in the development of chemistry than they once did.
8. Scientific progress and social changes are interrelated.
OXYGEN 41
| USING WHAT YOU HAVE LEARNED
Group A
1. Make as large a list as you can of substances containing
oxygen combined with other elements.
2. From what chemical compound was O2 first prepared
by Priestley?
3. From what compounds is O2 usually prepared in the
laboratory?
4. (a) What is a catalyst? (b) Illustrate your answer.
5. What assurance have we that the MnO., used in the lab-
oratory preparation of O2 acts as a catalytic agent?
6. What happens to the KCl that is formed by the de-
composition of KC1O3?
7. Why is it possible to collect O2 by the displacement of
water?
8. Priestley for a time collected gases by the displacement
of mercury. Why do we not use this method in collecting O.,?
9. List five physical properties of O.,.
10. Discuss the most important chemical property of O2.
11. (a) What is slow oxidation? (b) rapid oxidation? (c) Il-
lustrate each.
12. What is the kindling temperature of a substance?
13. Devise an experiment to show that air is necessary for
burning.
14. Name four substances whose kindling temperatures are
lower than that of coal.
15. Why is air removed from an electric-light bulb?
16. It is not as easy to burn a 500-page book whole as it is to
burn the same book a page at a time. Explain.
17. Why is asbestos used in making theater curtains?
18. (a) What are the conditions necessary for spontaneous
combustion? (b) Show how these conditions work out in the
spontaneous combustion of (1) finely divided coal in the en-
closed hold oi a ship, and (2) moist hay in a hayloft.
19. Why may the presence of dust in the air of a Hour mill
cause a frightful explosion?
20. Make a list of the uses of Or
21. Distinguish between combustion and burning.
(2 NEW WORLD OF CHEMISTRY
22. What is an oxidizing agent?
23. Describe the greatest industrial use of O2.
24. O2 rusts steel. What part does O2 play in derusting steel?
25. How do oxygen tents aid patients suffering from pneu-
monia and other respiratory diseases?
26. (a) What is basal metabolic rate? (b) How is it de-
termined? (c) What is its significance?
27. How does O2 play an important part in aviation?
. . T . . .
- I - • •
28. State the two most important industrial .sources of O0.
29. What three factors must be considered in selecting the
method for preparing large quantities of a substance?
30. (a) Describe a test tor O2. (b) What does the word test
mean to a chemist?
31. O2 and O3 can be changed easily from one to the other.
(a) How? (b) What great law does this illustrate?
32. How is O3 prepared?
33. How does O8 differ from O2 in physical properties?
34. By means ot an illustration, explain the meaning of
a I lot ropy.
35. What are the chief uses of O3?
36. Write word-equations for the burning of (a) coal (car-
bon) , (b) iron, (c) phosphorus, and (d) sultur.
Group B
37. By using mercury Priestley was able to collect a number
of gases that had escaped the attention of other scientists.
Explain.
38. Science today depends less upon chance discoveries than
it has in the past. Explain.
39. Window curtains behind a fish bowl filled with water
caught fire. Was this spontaneous combustion? Explain.
40. Why is O, passed through the inner tube ot the oxyhy-
drogen torch rather than through the outer tube?
41. Because ot oxidation, linseed oil hardens when exposed
to air. How is this chemical reaction speeded up?
42. Fishes die in air, man drowns in the sea. Explain.
43. Blow against a burning candle and it goes out. Blow on
the slowly dying embers of a fire and they burn more actively.
Explain.
44. Why is green, or moist, hay more susceptible to spon-
taneous combustion than dry hay?
OXYGEN 43
45. Illustrate the statement "progress in science and changes
in society are closely interrelated" by an example not men-
tioned in this chapter.
46. Coal dust is sometimes shoveled into a burning coal
furnace. Why?
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Have you ever heard of pure oxygen being administered
to athletes before a strenuous game? Consult the football or
track coach in your school or in a nearby college. What do
you conclude? Explain.
2. Organize a small group in your chemistry class and with
the help of your teacher set up apparatus such as Priestley first
used when he obtained oxygen. A large burning glass can be
borrowed from the physics department. Demonstrate the ex-
periment before your class.
3. Watch your mother put up fruit and jam in jars. De-
scribe her procedure, and explain why the food is heated,
and why a layer of paraffin is placed over the food before the
jar is sealed.
.4. Consult your doctor and a good book on first aid, and
write a report on the diagnosis and treatment of first-, second-,
and third-degree burns.
5. Read the story of a recent scientific discovery, (a) From
the account you read, what part do you conclude that chance
plays in scientific discoveries? (b) Why may the scientist
himself believe chance plays a more important part than it
actually does?
H Y D R O G E
LIGHTEST OF THE ELEMENTS
. . . Cavendish was almost passion-
less. An intellectual head thinking,
a pair of wonderful acute eyes ob-
serving, and a pair of skilful hands
* •. experimenting or recording are all
that I realize in reading his memoirs.
Cavendish did not stand aloof from
other men in proud spirit, he did
> , , so conscious of his inferiority, not
boasting of his excellence.
Dr. George Wilson, 1851
An eccentric man of science. While Priestley was performing his
immortal experiments, another Englishman was puttering around
in his palace laboratory on the wandering trail of phlogiston. This
man was Henry Cavendish, one of the most eccentric persons in the
whole history of science. It was said that he was "the richest among
the learned and the most learned among the rich." Although the
richest man in all England, he shut himself up in his private labora-
tory and spent more than 60 years tracking down many of the secrets
of nature.
In addition to his discoveries in chemistry, Cavendish was inter-
ested in physical problems. His work in the fields of heat and elec-
tricity was of highest rank; later it was followed up by the work of
other eminent scientists such as Joseph Black and Michael Faraday.
The elusive trail of phlogiston. Two hundred and fifty years
before the work of Cavendish, Paracelsus (par-d-sel'sus) of Switzer-
land had noticed bubbles of gas rise from an acid into which iron is
dropped. Although he found that this gas burns, he carried his
investigations no further. Then came Cavendish to whom the search
for truth was the ruling motive of life. He too noticed the gas
evolved when zinc or iron is dropped into an acid and went to work
HYDROGEN
45
to investigate this phenomenon. He collected the gas carefully, and
made a thorough study of it. He named it inflammable air because it
burned. He thought he had obtained phlogiston itself.
Strangely enough, the discovery of this gas, coupled with the dis-
covery of oxygen, paved the way for the complete overthrow of the
phlogiston theory and the establishment ol l.axoisier's true explana-
tion of burning. Though Priestley died still believing in the phlo-
giston theory, Cavendish, when the discussion over Lavoisier's new
chemistry became very heated, gave up his active interest in chemistry
and turned to the problem of determining the weight of the earth.
He said that he had no patience with squabbles, that he was inter-
ested in experimentation, not controversy.
How is hydrogen produced in the laboratory? Cavendish's lab-
oratory method of preparing hydrogen is still used. Zinc is placed
in a generator, as shown in Fig. 6, and dilute hydrochloric or sul-
turic acid is poured over it through a thistle tube. Bubbles of hydro-
gen gas form at once, and heat is generated. The hydrogen is col-
lected in the same way as oxygen, that is, by the displacement of
water. Why is hydrogen not collected by the displacement of air?
The chemical change that takes place is represented as follows:
Zinc -4- hydrochloric acid •
Zn + 2HC1
> zinc chloride 4- hydrogen
ZnCl2 4- H2
Hydrochloric acid is a compound of hydrogen and chlorine. The
zinc takes the place of, or replaces, the hydrogen of the acid and
liberates it as a free gas. Instead of hydrochloric acid, we now have
/inc chloride, which remains dissolved in the water in the generator.
Zinc chloride is a white solid.
All acids contain hydrogen, which may be replaced by certain
metals. Hence, to prepare hydrogen, we may use almost any acid and
one of a number of other metals instead of zinc. It is a curious fact
that when pure zinc is added to pure hydrochloric acid, the chemical
thistle tube
generator
bottle
Zn + HCI
delivery
Fig. 6. Laboratory prepa-
ration of hydrogen. Why
does thistle tube extend
below surface of liquid
in the generator?
hydrogen gas
collecting jar
- trough of water
46
NEW WORLD OF CHEMISTRY
„
Fig. 7. Preparation of hydrogen
by the action of sodium on water.
Perform carefully. The reaction
may be violent.
action is very slow. However, slightly impure zinc replaces the
hydrogen of the acid rapidly. The impurity in the zinc acts as a
catalyst.
Can hydrogen also be prepared from water? Perhaps you have
been thinking, "II the hydrogen of an acid can be replaced by a
metal, can the hydrogen of water also be replaced by a metal?" The
answer is Yes. Very active metals, such as sodium, potassium, and
calcium, possess enough chemical energy to replace the hydrogen of
water. This experiment may be tried by filling a bottle with water
and inverting it in a dish partly filled with water, as shown in the
illustration (Fig. 7) . Wrap in filter paper a piece of sodium the
si/.e of a pea and put the paper in the coiled end of a wire. Then
quickly insert this end of the wire into the bottle. Immediately, a
very active evolution of gas is noticed and the bottle becomes filled
with a colorless gas. The equation tor this reaction is:
Sodium -f water — » sodium hydroxide -f hydrogen
2Na +2HOH-> 2NaOH + H2
(Water may be represented as H2O or HOH.)
In the chemical reaction that occurs, sodium replaces half of the
hydrogen in water, and forms sodium hydroxide, NaOH. (The OH
group is called the hydroxyl group.) The sodium hydroxide remains
dissolved in the water. The water solution of sodium hydroxide feels
soapy, and turns pink litmus, a vegetable coloring matter, blue. By
evaporating the solution, the sodium hydroxide can be separated
from the water as a white solid. Sodium hydroxide belongs to a group
of compounds called bases (Chapter 14) .
Even iron, which is not nearly as active as sodium, will replace the
hydrogen of water while the water (as steam) is passed over red-hot
HYDROGEN
47
iron in a heated tube. In tact, this method ol prcp.ning hvdrogen
has been used to a slight extent, the products ol ihe ic.ution being
iron oxide and hydrogen.
What is the industrial method ol preparing hydrogen? In the
order ol the quantities ot gas produced, hydrogen is obtained lot
commercial use: (1) by passing steam oxer hot carbon, (2) by pass
ing steam through natural gas (methane) , in the presence ol a
catalyst, (3) by the electrolysis ol \\ater. The equations tor these
methods are:
1) Steam -f carbon — > carbon monoxide -f hydrogen
H«'9 ! !+."".".". ?.: -» CO + Hj
2) Steam -f methane — » carbon monoxide -f hydrogen
H>O + CH4 -* CO + 3H,
3 "i Water — > hydrogen -f oxygen
2H»O-> 2H2 4- O>
In the first two methods, the hydrogen mav be sepaiated from
tlie carbon monoxide gas by chilling the mixtuie ol gases. 1 lie cai-
bon monoxide, whose free/ing point is much higher than that of
hydrogen, solidifies, leaving nearly pure hydrogen gas behind. Only
a minor part ol the total hydrogen production is by electrolysis.
Physical properties of hydrogen. Hydrogen resembles oxygen in
most ol its physical characteristics. It is a colorless, odorless gas,
slightly soluble in water, and very difhcult to liquely. since it changes
Irom a gas to a liquid at — 252°C. It differs physically from oxygen
chiefly in its weight. It is the lightest element known. It is y1^ as
heavy as oxygen and ^^ as heavy as air. One liter ot hydrogen
weighs approximately 0.09 gram at standard conditions' of tempera-
ture and pressure.
The metals palladium and platinum absoib large volumes ol
hydrogen gas
porous cup
Fig. 8. Passage (diffusion) of
hydrogen through a porous
cup. Why is water forced
through the glass tube?
glass tube
water
48
NEW WORLD OF CHEMISTRY
hydrogen. This absorption, or occlusion, is accompanied by such an
increase in temperature that the metals actually glow. The absorp-
tion of hydrogen by these two metals takes place in one type of
lighting apparatus sometimes used in lighting the burners of gas
stoves. Such a lighter usually consists of a fine wire or wires of one
of the two metals strung between two suspension points. The lighter
is placed above the burner. When the gas is turned on, hydrogen is
absorbed by the wire. This causes the wire to give off heat. In about
a second, the kindling temperature of the gas is reached, and the gas
is ignited.
If two or more gases are at the same temperature, the particles of
the lighter gas move more rapidly than the particles of the heavier
gases. Since hydrogen is the lightest gas known, its particles are in
very rapid motion, passing through porous substances rapidly, as
shown in Fig. 8.
Hydrogen may burn quietly or explode violently. Although pure
hydrogen burns quietly in air or in oxygen with a pale blue, almost
colorless, flame, a mixture of hydrogen and oxygen may unite with
explosive violence. The two gases, when mixed and kept below a
temperature of about 800°C., will not unite; but a spark, a flame, or
a temperature of above 800°C. will cause them to unite violently. For
this reason, great care must be taken while experimenting with
hydrogen to keep all flames away from the generator. It is also nec-
essary to wait until the air has been completely expelled from the
generator^ In -lore setting fire to the hydrogen as it escapes from the
delivery wibe. ,
When the Frenchman, Pilatre de Rozier, heard of this gas which
Cavendish had studied, he tried an unusual and foolish experiment.
Fig. 9. Oxidation of hydrogen to form water. Will water form
if bell jar becomes hot? How may bell jar be kept cool?
HYDROGEN
49
He inhaled the gas until he had filled his lungs, and then as the
gas issued from his mouth he set (ire to it. All Paris held its sides
with laughter as it watched him spitting lire. However, when he set
fire in the same way to a mixture of this gas and air. "the conse-
quence was an explosion so dreadful that he imagined his teeth were
all blown out."
The chemical union of hydrogen and oxygen may be written:
2H2 -I- O2 -> 2H2O
This is an example of the strange behavior of chemical elements.
Hydrogen, a highly flammable gas, unites with oxygen, a gas which
helps things burn, forming water, a liquid that is one of the greatest
enemies of fire.
Experimental proof that water forms. Of course, when hydrogen
combines with oxygen, we do not see the formation of a flood of
water. We do not because the water that is formed at the tempera-
ture of burning hydrogen is invisible since it is in the form of water
vapor. However, it is possible to show the actual formation of water
by arranging the apparatus shown in the illustration (Fig. 9) . As
the invisible water vapor strikes the cool surface of the jar, drops of
a liquid form. This liquid is pure synthetic water. A synthetic com-
pound is a compound built up from simpler substances.
Hydrogen is a powerful reducing agent. Since hydrogen has
such a strong attraction, or affinity, for oxygen, it is able to tear
oxygen away from the other elements of many of its compounds.
The removal of oxygen from a compound is a cheniv.tl process
known as reduction. This ability of hydrogen to remove pxygen
from the other elements of a compound makes it a reducing or
deoxidizing agent.
Fig. 10. Reduction of copper oxide by hydrogen.
The fishtail burner is used to spread the flame,
thus heating a larger surface of copper oxide.
What is the function of the drying tube? Why is
water a product?
hydrogen generator
f*
50
NEW WORLD OF CHEMISTRY
For example, if pure hydrogen is passed over black copper oxide
brought to a red heat, as shown in Fig. 10, the hydrogen takes the
oxygen away from the copper oxide, leaving copper. This change
may be represented as follows:
Copper oxide (black) + hydrogen
CuO 4- H2
> water 4- copper (red)
» H2O -f Cu
Although hydrogen is a good reducing agent and is used to reduce
the oxides of such metals as wolfram and molybdenum, other re-
ducing agents are of greater commercial use. Perhaps the best
example of these is carbon (coke) , which is used in reducing iron
ore (iron oxide) to iron.
The relation of oxidation to reduction. In the experiment just1
described, the copper oxide, CuO, is reduced to copper. At the same
time hydrogen is oxidi/ed to water, H2O. This illustrates a general
principle: namely, that the reducing agent is itself always oxidi/ed.
You should remember that whenever one substance is reduced,
another is oxidi/.ed. Thus oxidation and reduction always occur in
the same reaction.
The bunsen burner and how it works. One of the most familiar
pieces of apparatus in the chemical laboratory is the bunsen burner
(named after Robert Bunsen, a German chemist who introduced it in
1855) . The function of this burner is to mix a gaseous fuel with
air in order to make a lumliiminous flame that has a high tempera-
ture and will not deposit soot.
The parts of the bunsen burner designed for burning manufac-
tured gas are shown in Fig! 1 1. Gas enters through the side of the
stand, and its speed is increased by passing through the narrow spud.
The rapidly moving gas draws in air through the hole in the collar,
which may be adjusted to permit the correct volume of air to enter.
The mixture of gas and air passes up through the barrel and is
ignited at the top.
When the pressure of the gas is low, the flame may be drawn back
and may burn at the spud. This striking back of the flame can be
Fig. II.
of m
•r.
of
Item**
n
of
of
of
combustion)
• It*
HYDROGEN
51
dangerous because the stand -of the burner becomes hot. enough to
produce serious burns when touched, and occasionally the hot burner
melts the rubber hose and ignites the gas.
In the type of burner for natural gas, the air intake is larger
and the hole in the spud is much smaller. It is also supplied with
a flame retainer which has 0 small pilot jets led by gas bled from
the main tube. This eliminates the tendency tor the flame to blow
out.
The burner on a gas range consists of a series of small bunsen
burners. Not only is the nonluminous ilamc of the bunsen burner
used in cooking, but it is used also to some extent in illumination by
burning the gas inside gas mantles consisting of 09 pen cut thorium
oxide and 1 percent cerium oxide. These oxides, when heated, glow
with a rich white flame. They were first made commercially practical
by von Welsbach (ton vels'baiO . The widespread use of electricity
as an illimiinam has made this type of mantle burner all but obso-
lete.
The nature of a flame. A flame is produced only when combustible
vapors reach their kindling temperature in the presence of air,
oxygen, or some other substance that supports combustion. When
iron is heated gently, it glows, but a flame is not produced, because
iron does not vapori/e at low temperatures. A candle, on the other
hand, burns with a flame, because the candle melts and the heat
from the burning wick causes this licjuid to change to a vapor that
burns in air. It can be shown that the interior of a candle flame
consists of a combustible gas by holding one end of a piping hot
tube in the innermost part of the flame and then touching the other
end with a lighted match (Fig. 12) .
The structure of the bunsen flame. Three distinct /ones are no-
ticeable in the bunsen burner flame as well as in the candle flame.
The innermost zone A is composed of combustible gas that has not
yet reached its kindling temperature. This fact may be tested by
placing a match head in zone //. This may be clone by piercing
a match near the head with a pin. The pin serves as a bridge across
burning gas
Fig. 12. Experiment thawing
j the nature of the innermost
: zone of a candle flame.
52
NEW WORLD OF CHEMISTRY
the opening of the burner. If the burner is lighted carefully, the
matcli will not catch fire. Outside zone A is zone B in which the gas
is burning. The light-purple zone C is the region of complete com-
bustion, where carbon dioxide and water vapor are formed.
Just below the tip of the flame, plenty of air is available and,
therefore, this is the oxidizing part of the flame. Zone B is the region
of somewhat incomplete combustion, where one of the products
formed is carbon monoxide, a reducing agent. Hence this zone is
the reducing part of the flame. This zone of the flame is used in
reducing metallic oxides to free metals by means of carbon. The
blowpipe directs the flame, as shown in Fig. 71, page 321.
The luminosity of the flame of the bunsen burner, when the collar
is closed, is caused by the decomposition by heat of a small amount
of the hydrocarbons and the subsequent burning of free carbon. This
may be proved by holding a cold dish in the luminous flame. The
carbon is reduced below its kindling temperature by contact with the
dish and, hence, deposits soot on the dish.
Hydrogen rides the winds. Interesting historically, but of only
slight commercial importance today, is the use of hydrogen in filling
balloons and other lighter-than-air craft. Soon after the discovery
of hydrogen. Dr. Charles (shiirl) , of Paris, constructed the first
large hydrogen-filled balloon, and in the presence of 300,000 spec-
tators, Pilatre de Ro/ier, who had experimented unwisely with
hydrogen before, bravely climbed inside its basket and started on
the first aerial voyage ever made by a human being. Since that time,
balloons filled with hydrogen have carried men around the world
and have lifted explorers of the atmosphere to altitudes of almost
14 miles. The development of the motor-propelled, rigid balloon
called a dirigible gave us gigantic craft, nearly 1000 feet long and
weighing more than 100 tons.
From "Photographic History of the Civil War" by Albert Shaw
The observation balloon
was used to advantage by
the Union Army in the War
Between the States. Here
the balloon Intrepid is being
filled with hydrogen by the
generators at the left.
Such helium-filled bal-
loons, carrying recording!
• instruments, are used to
determine weather condi-
tions at high altitudes.
Official Unttcd States A'nrj/ J'/mf,, r ,-,,,,-,
The history of this type of flying machine was filled with tragedy,
to a large extent because hydrogen ignites and explodes readily.
With the discovery of large sources of a nonflammable gas, helium,
and its subsequent use, this ever-present danger was removed. There
are no dirigibles in existence today, although blimps, which are
motor-propelled, nonrigid balloons, are used for military purposes.
Helium, which possesses about 93 percent of the lilting power of
hydrogen, may be obtained today at a cost not much greater than
that of hydrogen. The United States is fortunate in possessing the
largest supply of helium of any nation. As high as seven percent
helium is extracted from the natural gas of some western fields.
To conserve our own supplies and prevent other nations from
using helium for military purposes, Congress passed a law in 1938
which placed its export under very strict government control. Sci-
entific discoveries often turn out to be two-edged swords, and the
careful control of such discoveries should be part of the business
of government.
Hydrogen, a gas, helps solidify fats and oils. One of the most
important uses of hydrogen is based on the discovery that hydrogen
Procter and Gamble Compan\
These machines, called
"freezers," are used in the
hydrogenation of vege-
table oils to give the oils
creaminess and smooth
texture.
54
NEW WORLD OF CHEMISTRY
can be chemically combined with other substances to form new
products of great value. This process of chemically combining hydro-
gen with other substances is called hydrogenation. In normal times,
hydrogenaiion ol fats and oils is the largest use of hydrogen.
Generally, during hydrogenation, liquid oils are changed to semi-
solid fats. Thus, cottonseed oil, in the presence of a finely divided
nickel catalyst, unites with hydrogen, forming a white fat that is
often sold nuclei such trademarks as "Crisco" or "Spry." Fish oils,
formerly almost useless, have been hydrogenated similarly and ren-
dered useful. Fats for soap-making and candle-making have been
prepared commercially by this process also. Crude oil, coal, and cer-
tain waste products of petroleum refining are also hydrogenated and
changed in this way into high-quality gasoline. Wood alcohol is an-
other important chemical made with the aid of hydrogen.
Oleomargarine, or margarine, a widely used butter substitute, is
made of either hydrogenated vegetable oils or animal fats, or mix-
tures of these substances. Much of the margarine now marketed is of
vegetable origin.
The introduction of hydrogenation on a large scale helped the
cotton farmers of the South by opening a new market for their cot-
tonseed oil. It has had even greater effects abroad, for it has enabled
several Kuropean countries to become less dependent upon other
nations for their vital supplies of gasoline and edible fats.
Hydrogen for the synthesis of ammonia. Ammonia, a compound
of nitrogen and hydrogen, leads a double life. It is an important
component of nearly all of the world's most powerful explosives,
and it is an important ingredient in fertili/ers used all over the
world. The most significant industrial process for the preparation of
ammonia is one that embodies the direct combination, or synthesis,
of nitrogen and hydrogen. This process is discussed in detail in
Chapter 17. In wartime, production of ammonia probably requires
more hydrogen than any other industrial use.
In this plant, bituminous coal it
hydrogenated to produce many
useful chemicals.
An illustration of a solar prominence. Solar prominences are great tongues of
glowing hydrogen which shoot out of the chromosphere of the sun and extend
far into space. The flames often attain lengths of more than 100,000 miles and
have been known to reach lengths of more than 1,000,000 miles.
Hydrogen in heating and cooling. Hydrogen has another impor-
tant use. When it is burned, it gives off four times as much heat as
an equal weight of coal. Because of the high temperature produced,
hydrogen is used both in the oxy hydrogen torch and as a constituent
of certain gaseous fuels such as water grw, which contains about 50
percent hydrogen. Hydrogen gas is also used as a cooling agent in-
stead of air. Because of its very low density it cuts down friction
in machines such as high speed turbine-generators. When hydrogen
is so used, the machine is completely enclosed to shut out air.
The test for hydrogen. Hydrogen is not the only colorless gas that
burns with a pale blue, almost invisible flame. But it is the only gas
that forms water as the only product of its burning. This fact gives
us a simple chemical test for hydrogen.
Where is hydrogen found? Unlike oxygen, only small amounts of
hydrogen occur on earth in the free state. However, live hydrogen
is the most abundant element in the sun. The immense luminous
tongues, or solar prominences, some of which extend half a million
miles from the sun's surface, consist of glowing hydrogen. It is also
the commonest element found in interstellar space, and by far the
most abundant material out of which the whole universe is built.
Combined hydrogen is very common on earth. Hydrogen consti-
tutes about 11 percent by weight of all water and is one of the ele-
ments in petroleum, all acids, and living cells (protoplasm). In
spite of its widespread occurrence, the extreme lightness of hydrogen
accounts for the fact that it constitutes only 1 percent by weight of
the earth.
55
Henry Cavendish (1731-1810), one of
the most unusual personalities in the
history of science. For the most part, he
shunned society of any kind, even in-
structing his servants to keep out of his
sight. Yet, over his long lifetime, he
made many important contributions in
the fields of chemistry and physics.
Theories lead to great discoveries. In 1932, three Americans
headed by Harold C. Urey discovered that ordinary hydrogen could
be separated into two distinct forms. They named the heavier of
these forms deuterium. Two years later, it was proved that there
is a third form of hydrogen. This form, the heaviest of the three,
has been named tritium. The three forms have the same chemical
properties, but differ in certain physical properties.
This achievement is remarkable, not so much because ordinary
hydrogen was shown to be a mixture of three forms of the same
element, but because it afforded a definite example of a great con-
tribution to chemistry made by scientists who forecast these dis-
coveries purely from theory. Their success points to the fact that
science needs the man who experiments, the thinker who can put
his imagination and reason to work in propounding theories, and the
engineer who works to discover how these theories and processes
may be used in the service of man.
Another example of how theoretical problems in science may turn
out to be of great practical value, was the theorizing of the great
mathematician, James Clerk Maxwell. In 1863 he came to the con-
clusion that just as light results from a wave disturbance in the
ether, so electric disturbances from a spark should produce similar
waves, invisible, to be sure, but nevertheless existent. Experimental
evidence of such waves was found 23 years later by a young physicist,
Heinrich Hert/. These Hertzian waves, now known as radio waves,
were later used by Marconi in the transmission of wireless messages.
Thus modern radio originated, and out of a purely theoretical in-
vestigation came one of the most practical and valuable of modern
scientific marvels. Often, even the greatest scientist cannot predict
the practical value of theoretical research.
56
HYDROGEN
57
YOU WILL ENJOY READING
Hoyle, Fred. The Nature of the Universe. Harper and
Brothers, New York, 1950. This is a very small book that gives
an exciting picture of the composition of the stars and inter-
stellar space.
Jaffe, Bernard. Men of Science in America, pp. 3S9-355.
Simon & Schuster, New York, 1944. Interesting inioimation on
early American work in aeronautics.
Walters, Leslie. "Chemistry Exhibits and Projects." Jour-
nal of Chemical Education, March, 1939, pp. 113-115. An
illustrated article on exhibits and projects made by high-
school students. Suggestions for what you, too, could do along
this line.
USEFUL IDEAS DEVELOPED
1. A reducing agent removes oxygen from a compound con-
taining oxygen. (This definition will be somewhat expanded
later.)
2. Reduction and oxidation always occur in the- s.nnc re-
action.
3. Occlusion is the absorption of gases by metals.
4. Careful control of some scientific discoveries is necessary
to prevent their misuse.
5. Chemistry needs the experimental chemist, the scientist
able to use creative imagination in formulating theories, and
the engineer who works to discover how theories and processes
may be used in the service of man.
USING WHAT YOU HAVE LEARNED
Group A
1. (a) Who is credited with the discovery of pure H2?
(b) How was H2 first prepared?
2. (a) What substances are used in the laboratory prep-
aration of H2? (b) How do we know that the H.. comes iroin
the acid and not from the metal? (c) Write the word-equation
for this reaction.
3. (a) What advantage is there in collecting H2 by the
displacement of H2O? (b) What kind of gas could not be
collected in this manner?
58 NEW WORLD OF CHEMISTRY
4. (a) Make a labeled diagram of the laboratory prepara-
tion of H2. (b) Why should the thistle tube extend below the
surface of the liquid in the generating bottle? (c) What are
two reasons for using a thistle tube?
5. (a) When Na displaces H2 from H2O, what is the other
substance formed? (b) Why can you not see it? (c) How can
you obtain it for inspection of its properties?
6. What metals other than Na are so active that they will
liberate H2 from water when the metal is simply placed on
H2O?
7. (a) What metal will liberate H2 from H2O under cer-
tain special conditions? (b) Describe this method of prepar-
ing H2.
8. (a) In what two ways is HL, prepared tor commercial
use? (b) Why were these methods selected in preference to
ones used in the laboratory?
9. What precaution must be observed in preparing and
handling H2?
10. (a) List five physical properties of H2. (b) Describe a
simple experiment to illustrate the tact that H2 is a very light
gas.
11. Determine the weight and cost oi the H2 that would be
needed to fill a dirigible of seven-million-cubic-foot capacity.
Consider that the cost of the H2 is $2.00 per hundred cubic
feet of gas.
12. What element in the air is used in the burning of H2?
w
13. How can you show that water is formed when hydrogen
burns?
14. What is reduction?
15. In the reduction of CuO by H2, what substance is oxi-
dized?
16. How can you identify each of the products obtained in
the reduction of CuO by H2?
17. Oxidation and reduction always occur together. Ex-
plain.
18. (a) Give an example of reduction carried out on a large
scale in industry, (b) What reducing agent is used?
HYDROGEN 59
19. (a) What is the function of each part of a bunsen
burner? (b) What is meant by the striking back of the flame?
20. (a) What conditions are necessary for a flame? (b)
Make a labeled drawing of the flame of a bunsen burner.
21. What formerly useless byproduct is converted into a
very useful substance by the use of H0?
22. How is oleomargarine made?
23. In what special fuels is H2 chiefly responsible for the
high temperatures obtained?
Group B
24. Can you suggest a sale way to test H, for its burnability
as it emerges from the end of the delivery tube in the labora-
tory preparation?
25. Helium gas is twice as dense as hydrogen gas and yet
can lift about 93% as much weight as hydrogen gas. Explain.
26. (a) Could pure H2 be used in the gas range at home?
(b) Explain.
27. Why does pure H2 burn quietly in an atmosphere of air,
yet burn with explosive violence when the two gases are mixed
and ignited?
28. Tightly bound inflated balloons gradually collapse. Ex-
plain.
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. A student's mother stopped using "Crisco" after her son
came home with the news that "Crisco" was not a "natural
product" but was manufactured chemically, (a) Was the
mother justified? (b) Give reasons for your answer after
you have consulted your family physician, the producers of
"Crisco," and the U.S. Department of Agriculture.
2. Organize a small discussion group in your chemistry class
and discuss the topic, "Synthetic chemistry has helped in the
rise of totalitarian states."
3. Prepare a report or organi/e a class discussion on the
topic "Theoretical science has irequently resulted in great
practical discoveries."
4. WATER:
THE COMMONEST LIQUID COMPOUND
. . . Laboratories are necessary, and,
though an artist without a studio or
an evangelist without a church might
conceivably find under the blue
dome of heaven a substitute, a sci-
entific man without a laboratory is a
misnomer. Frederick Soddy, 1877-
Water a compound of two gases — impossible! For thousands of
years, water was considered an element. Aristotle, one of the wisest
of Greeks, included this liquid among the "four elements" of the
ancients. No power of man seemed strong enough to break it up
into any recognizable components. Although it is true that by 1780
a number of scientists had really decomposed water so that hydrogen
was liberated, they were unaware of the real nature of what they
had done. They could not believe that hydrogen had actually come
from the water.
"It is very extraordinary that this fact should have hitherto been
overlooked by chemists. Indeed, it strongly proves that in chemistry
it is extremely difficult to overcome prejudices imbibed in early
education." These were the words of Lavoisier, Often it is hard to
overcome superstition, prejudice, and tradition. However, we must
learn to do so in order to think and act scientifically.
In 1784, Cavendish, who had studied hydrogen, read an exciting
paper before the members of the Royal Society of England. This
is what he told them: "Water is a compound of oxygen and hydro-
gen." What a startling announcement! Water a compound of two
colorless, tasteless gases? What were his proofs? Cavendish told them
60
WATER
61
quietly and without emotion. He said that he had placed in a glass
flask a mixture composed of about twice as much air as hydrogen.
Then he had passed an electric spark through the mixture. "All the
hydrogen and about one-fifth of the air condensed into a dew which
lined the glass. In short," he continued, "it seemed pure water." His
experiments proved conclusively that water is a compound of oxygen
and hydrogen, and yet Cavendish said "it seemed." He suspected his
listeners would not be convinced. Water a compound of two gases —
impossible!
Lavoisier convinces the world that water is a compound. Lavoisier,
who had explained the nature of burning, determined to tear apart,
or analyze, water by an experiment that would convince the world,
just as Cavendish had shown the world that he had built up, or
synthesized, water from oxygen and hydrogen.
He arranged the apparatus shown below. In the retort A he heated
pure water, so that steam would pass through the tube containing
pure charcoal, which, as you see, was heated in a furnace. The gas
that escaped passed into the jar H. He found carbon dioxide gas
dissolved in the water in the jar and identified the gas issuing from
H as hydrogen. The water (steam) had been broken up into hydro-
gen, which passed on as a gas, and oxygen, which combined with the
carbon, forming carbon dioxide. Lavoisier collected and weighed
both gases. The weight of all the resulting products accounted for
all changes in the weights of the substances used in the reaction.
Thus Lavoisier proved without a doubt that water could be broken
up into the two gases of which it is composed — hydrogen and
oxygen.
Fig. 13. Apparatus used by Lavoisier to
analyze water. Compare this equipment with
that used in a modern laboratory.
burning charcoal
water
stopcock
hydrogen gas
reservoir of water containing
some H2SO4
platinum
cathode
Fig. 14. Electrolysis of water. Why is a reservoir necessary in this apparatus?
The equation for the chemical changes that took place in this
experiment is:
2H2O + C -> 2H2 + CO2 (carbon dioxide)
How water can be broken up by an electric current. In 1800, the
invention of the electric battery by Volta put into the hands of
scientists a new and powerful tool for determining the composition
of compounds. The decomposition of any compound by electricity
is called electrolysis.
Within a few months following Volta's invention, an apparatus
had been devised for the electrolysis of water. This special piece of
apparatus is used in the laboratory in determining the composition
of water by volume. Water is poured into the center glass tube,
which acts as a reservoir, and fills the other two tubes, as shown
in Fig. 14. Since pure water is an extremely poor conductor of
electricity, that is, will not allow electricity to pass through it readily,
a small amount of sulfuric acid is added. (At the close of the chem-
ical action, the amount of sulfuric acid is unchanged.)
The electric current enters the apparatus through one of the
platinum electrodes, passes through the mixture of water and sulfuric
acid, and leaves the apparatus at the second platinum electrode.
The source of current may be any source of direct current, such as
a storage battery or dry cells. As soon as the circuit is closed, bubbles
of gas collect on the surface of the electrodes, rise through the water,
and collect at the top of the two outside tubes. The gas that collects
at the negative electrode, or cathode, is hydrogen; the gas that col-
lects at the positive electrode, or anode, is oxygen.
WATER
63
The composition of water by volume. No matter when we stop
the action of the current in the electrolysis of water, we discover a
singular fact. The volume of the gas at the cathode is twice the
volume of the gas at the anode. In other words, water always consists
of 2 parts of hydrogen to 1 part of oxygen by volume.
If the gases hydrogen and oxygen are mixed in a closed chamber
and then exploded by an electric spark, water vapor is formed, and
the ratio in which these gases unite, forming water vapor, is always
2 volumes of hydrogen to 1 of oxygen. Any excess of either gas is left
uncombined.
The composition of water by weight. If pure water is thus de-
composed, forming two gases in a definite ratio by volume (vol-
umetric ratio) , and if water is formed by combining these gases in
the same volumetric ratio, it ought to be possible to find in this com-
bination a constant ratio by weight. This is what Cavendish did in
1784 when he formed water by combining these two gases. He actu-
ally weighed both the gases and the water. Today we know that if
we combine 1 gram of hydrogen (11 liters) with 8 grams of oxygen
(5.5 liters) , we obtain 9 grams of water. Thus, the composition of
water by weight is 1 part of hydrogen to 8 parts of oxygen.
The composition of water by weight may be tested in the labora-
tory by making use of the fact that pure hydrogen passed over heated
copper oxide reduces the copper oxide, yielding free copper, by
uniting with the oxygen of the copper oxide, thus forming water
(see page 50) . The weight of the water formed is always found
to be exactly the same as the weight of the hydrogen used plus the
weight of the oxygen lost by the copper oxide. The following data
are the results of such an experiment:
Weight of copper oxide before experiment 80 grams
Weight of copper left after experiment 64 grams
Weight of oxygen that combined with hydrogen 16 grams
Weight of water produced 18 grams
Weight of hydrogen used 2 grams
Result: 16 parts by weight of oxygen combined with two parts of
hydrogen (or 8 parts of oxygen combined with one part of hydro-
gen) , forming water.
. induction
latmum electrodes co;|
Fig. 15. Eudiometer. In this appara-
tus, a mixture of hydrogen and oxy-
gen is ignited by an electric spark,
forming water. Any excess of either
gas acts as a cushion against the
rush of mercury up the tube.
Accuracy and satisfaction in scientific investigations. An exact
determination of the relative combining weights of oxygen and hy-
drogen was essential to the mathematics of chemistry. Hundreds of
men performed thousands of experiments to determine these values
as accurately as human ingenuity could devise. Many eminent Ameri-
cans were among them. Edward Morley (1838-1923) , a professor at
Western Reserve University, spent more than ten years of his life on
such experiments. Finally he arrived at a number which still stands
as the basis for chemical calculations. He found that 1.008 parts of
hydrogen by weight combine with eight parts of oxygen by weight,
forming water.
Morley and the others did not receive monetary rewards from
these experiments. Theirs was a labor of love, a work of pure scien-
tific research. Their sole reward was the satisfaction of knowing that
they were helping to increase scientific knowledge whjch might be
of use to all humanity. The world needs such unselfish men.
The law of definite proportions. You have learned that the com-
position of water is always the same. Perhaps you have wondered
if this is a property of all chemical compounds. It is one of the
fundamental laws of chemistry that, in forming compounds, elements
combine in exact proportions. In fact, constancy of composition is
the most valid test used in deciding whether a substance is a com-
pound or a mixture. In other words, the composition of a pure com-
pound never varies. This is the law of definite proportions..
"The stones and soil beneath our feet and the ponderous moun-
tains are not mere confused masses of matter; they are pervaded
throughout their innermost constitution by the harmony of num-
bers." This is indeed fortunate, for if the composition of pure com-
pounds ever changed, instead of being always the same, the exact
measurements that we use in chemistry would not be possible. Par-
ticularly is this true of quantitative chemistry, for if the composition
of compounds varied, we should have no reliable standard chem-
icals. Standards are necessary, not only as a means of judging the
purity of a substance which we plan to use, but also as a reliable
basis for comparison.
64
WATER 65
What is a "pure" chemical? Absolutely pure chemicals are almost
impossible to make or buy. Chemicals that have no appreciable trace
of impurities are called C.P., that is, chemically pure. The designa-
tion U.S.P. refers to standards of purity of chemicals to be used in
medicines listed in the United States Pharmacopoeia. Such chem-
icals contain no harmful impurities. Reagent grade chemicals con-
form to the standards of the American Chemical Society. Tech chem-
icals do not meet any definite, or fixed, standards of purity but are
suitable for many uses in which slight impurities are of little im-
portance.
Patent medicines, other remedies, and drugs, such as aspirin, are
frequently advertised and sold under trademarks. The purchaser
should always carefully examine the label on the package in order
to be sure of the purity of the contents. The chemical composition
of the contents will also serve as a guide to the fair price of the
article. Occasionally, simple and inexpensive C.P. chemicals, such
as bicarbonate of soda, are masked behind trademarks and sold at
exorbitant prices. The consumer can help to guard against such
practices by insisting that accurate and complete statements of com-
position be printed on labels of all packaged goods.
Some physical properties of water. As you know, water is an odor-
less, tasteless liquid that is colorless, except in very thick layers, when
it appears blue. Pure water freezes at 0°C. or 32°F. and boils at
100°C. or 212°F., at standard conditions. In general, impure water
has a higher boiling point and a lower freezing point than pure
water.
The fact that water is so universally distributed has led to its use
in the devising of scientific standards of measure. Thus a gram, the
metric unit of weight, is, by definition, the weight of a milliliter * of
chemically pure water at 4°C. This temperature is chosen because,
when water cools, it contracts until the temperature reaches 4°C.
Below that temperature it begins to expand again. Hence, 4°C. is
the temperature at which a unit volume of water weighs the most.
The weight of a unit volume of a substance* is known as its density.
Thus at 4°C. water has its greatest density. Since we use the gram
and the cubic centimeter or the milliliter as our units of weight and
volume, we may redefine the density of a substance as the weight in
grams of 1 cubic centimeter or 1 milliliter of that substance.
* A milliliter (ml.) is a small unit of capacity (volume) in the Metric System.
Numerically it is equal to 0.001 of a liter and is used in measuring the volume of
fluids. The cubic centimeter (cc.) is also used in measuring volumes, particularly
of solids.
66
NEW WORLD OF CHEMISTRY
0.9
0.25
cork
aluminum
water 1.0
Fig. 16. The relationship of specific gravity and
buoyancy. Aluminum has a specific gravity greater
than that of water and does not float. Ice has a specific
gravity slightly less than that of water and floats
largely submerged. Cork has a low specific gravity
and floats with most of its mass above water.
Since water has a density of 7, that is, 1 milliliter of water weighs
1 gram at 4°£., the density of any substance is also the ratio of the
weight of a given volume of that substance to the weight of an equal
volume of water. We call this ratio the specific gravity (sp. gr.) of
the substance. It shows the comparison between the weight of the
substance and the weight of an equal volume of water. For example,
when we say that concentrated su If uric acid has a specific gravity
of 1.84, we mean that it is 1.84 times as heavy as water, volume for
volume. Since below 4°C. water expands, ice is lighter than water
and floats on it. Ice, of course, has a specific gravity less than that of
wratcr.
TABLE 1. DENSITIES OF COMMON SUBSTANCES P
,.vvvvvvv>.vvvvv\.vvvvvvvvvvvvvvvvvvvv^^^
Sodium
Water
(In grams per cubic centimeter or per milliliter)
0.97 I Iron (pure) 7.86 | Lead
1
Gold
19.3 I Platinum
11.37
21.5
WATER 67
The high specific heat of water. The amount of heat necessary to
raise the temperature of 1 gram of water 1 degree centigrade is called
a calory. The number of calories necessary to raise the tempera-
ture of 1 gram of a substance 1 degree centigrade is known as the
specific heat (sp. lit.) of that substance. Because it takes 1 calory to
raise the temperature of 1 gram of water 1 degree centigrade, the
specific heat of water is 1.
Since it takes only one-thirtieth as much heat to raise the tempera-
ture of 1 gram of mercury 1 degree centigrade as it does to raise 1
gram of water 1 degree centigrade, the specific heat of mercury is one-
thirtieth of 1, or^j.
Water has a higher specific heat than most other substances. Since
it requires so much heat to raise its temperature, it warms up slowly.
Conversely, upon cooling it gives up a larger amount of heat for the
same fall in temperature than most other substances do. It is partly
because of the high specific heat of water that it is used in home heat-
ing systems and in the cooling systems of automobiles.*
The chemical properties of water. Water is a stable compound,
that is, it cannot be decomposed easily. It does not even begin to
decompose into hydrogen and oxygen until a temperature of
1000°C. is reached. Even at 2500°C. only two percent of it is decom-
posed. However, electricity, in the presence of a catalyst, tears it
apart easily (see page 62) .
At ordinary temperatures, water is decomposed by the more active
metals, such as sodium and potassium, and at higher temperatures
by the less active metals, such as iron. In these cases the gas liberated
is hydrogen. Water is also decomposed by the more active nonmetals,
such as chlorine and bromine, but these liberate oxygen from water
instead of hydrogen.
Water acts as a catalyst in many chemical reactions. For example,
perfectly dry oxygen and hydrogen do not unite when a spark is
passed through them, yet the faintest trace of water causes such a
mixture to explode. Phosphorus does not burn in perfectly dry air,
but burns readily if even a trace of water vapor is present.
* Although the temperatures of boiling water and steam are the same, it takes
about 540 calories to change 1 gram of water from its boiling point of 100°C.
to steam at 100°C. This amount of heat is called the heat of vaporization of
water. Real steam, or water vapor, the gaseous form of water, is invisible. The
visible cloud commonly called steam is water vapor after it has condensed into
tiny liquid droplets.
In ice water the temperatures of the freezing water and the melting ice are
both 0°C. Yet it requires about 80 calories to change 1 gram of ice at 0°C. to
1 gram of water at 0°C. This amount of heat is called the heat of fusion of ice.
(left) Photomicrograph of a crystal of sodium carbonate, (right) The same crys-
tal after a few hours of exposure to air. What has occurred?
Water of crystallization. A crystal is a solid mass having a well-
clefined and angular form. The word is derived from a Greek word
meaning clear ice. Most elements and compounds are capable of as-
suming the crystalline form, showing sharp edges and flat surfaces.
Such a substance is crystalline washing soda, or sodium carbonate.
When a crystal of washing soda is heated or even exposed to air,
it gives off water and crumbles to a white powder which is not
crystalline. The weight of water liberated bears a fixed ratio to the
weight of the crystal and is united chemically with the compound of
which the crystal is composed. Water which is thus chemically united
with a substance and gives that substance its crystalline form is called
ivater of crystallization. Such water is rather loosely held in chemical
combination and may be easily expelled. The water of Crystallization
is separated from the rest of the formula by a centered dot, which
means plus (+) and is not a multiplication sign. A substance that
contains water of crystallization is sometimes called a hydrate.
Another common hydrate is crystallized copper sulfate once com-
monly known as blue vitriol. When this compound is heated, its
water of crystallization is liberated and it crumbles to a white powder.
CuS04 5H20 -> CuS04 + 5H20
crystallized copper sulfate water of
copper sulfate (anhydrous) crystallization
This change in color is further evidence that the water of crystalliza-
tion is chemically united with the copper sulfate. Use is made of the
difference in color between white anhydrous copper sulfate and the
blue hydrated copper sulfate as a test for water. Water will change
anhydrous copper sulfate to the blue hydrate.
The ability to form crystals is not always dependent upon the pres-
ence of water. Many crystalline substances, such as table salt (so-
dium chloride) arid sugar, do not contain water of crystallization.
They are said to be anhydrous, meaning without water. Crystals that
have lost their water of crystallization are also said to be anhydrous.
68
WATER
69
Efflorescent substances give up water. Crystallized washing soda,
on exposure to air, loses its water of crystallization and crumbles
to a powder. Such a substance is said to be efflorescent, which means
that it gives up its water of crystallization on exposure to air. The
drier the air, the faster the loss of water of crystallization.
Deliquescent substances take up water. Dry sodium hydroxide,
when left exposed to air, soon absorbs enough water from the at-
mosphere to dissolve itself in this water. Such a substance is said to
be deliquescent. The higher the percentage of water vapor in the
air, the faster the process of deliquescence.
Calcium chloride, a white solid, is deliquescent and is often used
to sprinkle dry roads and tennis courts. It absorbs moisture from the
air and, in this way, helps to keep the dust down. Magnesium chlo-
ride, an impurity found in common table salt, is also deliquescent.
Removal of the magnesium chloride causes pure table salt to remain
dry in damp weather, to pour easily, and not to cake. (What sub-
sance does your mother put in a salt shaker to keep the salt from
becoming lumpy?)
Deliquescent substances may be used as drying, or dehydrating,
agents. For example, concentrated sulfuric acid absorbs moisture
from the air and, therefore, is used in drying gases. When used in
the laboratory, these dehydrating agents (the most common of which
are sodium hydroxide, sulfuric acid, and calcium chloride) are often
placed in the lower compartment of a vessel known as a desiccator;
the upper compartment, only partially separated from the lower,
contains the substance to be dried.
Our lives depend on water. Water is essential to life. Almost 70
percent of the total weight of the human body is water, and plants
contain even more. Lettuce, for example, contains as much as 95
percent water by weight.
,,.. .:-• *v!-'" -„ Fig. IF* Desiccator. 5yb»
i. * to be are
ln the VGpw «m4 the
is put In
of
70 NEW WORLD OF CHEMISTRY
Immense tracts of land in our own country, such as the hitherto
arid Columbia Basin of the Northwest, have been or will be turned
into rich farmland by irrigation. Federal government projects have
included construction of huge dams such as the Norris Dam, the
Hoover Dam, the Fort Peck Dam, and the Grand Coulee Dam. Be-
hind mountainous walls of earth and concrete are stored huge res-
ervoirs of water which are changing more wasteland into fertile
plains and are helping to solve the problem of the frequent recur-
rence of disastrous floods in certain areas.
Water is found in rocks, paper, fibers, and other substances gen-
erally thought of as "dry." The pages of this book may contain as
much as 10 percent water by weight. The importance of water, the
most common solvent in nature, is discussed in detail in Chapter 15.
Water and health. Water is a major component of all body tissues
and fluids. It plays a major role in the preparation of foods we eat
and in the processes of digestion and assimilation. Both nutrients and
oxygen are carried to the cells of the body in fluids composed chiefly
of water, and many of the waste products of the body are carried
away and eliminated in a similar mariner. \Vater, in the form of
perspiration, aids in regulating the temperature of the body.
To maintain normal body processes, rather large quantities of
water are necessary. The amount of water a person ^should drink
each day to enable these processes to be carried on varies with the
kind and amount of activity, the temperature, and various other
factors. However, 6 glasses each day should be considered a minimum
for good health.
The idea that water should not be drunk with meals is without
foundation. Digestion proceeds normally even when large quantities
of water are present in the stomach. However, it is very important
not to substitute the drinking of water with meals for proper and
complete chewing.
This illustration shows how a single calcium chloride crystal deliquesces, gradually
absorbing water from the air until it becomes completely dissolved.
WATER
71
YOU WILL ENJOY READING
Holmes, Harry N. Out of the Test Tube (4th ed.) ,
Chap. VII. Emerson Books, New York, 1945. "The Elixir of
Life," of course, refers to water.
Read the Label on Foods, Drugs, Devices, Cosmetics. Cata-
logue No. FS 13.111:3/2, 1953. Supt. of Documents, Govt.
Printing Office, Washington, D.C. Revision No. 1 of a 35-page
illustrated pamphlet containing valuable information. 15^
Thorpe, T. E. Essays on Historical Chemistry, pp. 98-122.
The Macmillan Co., London, 1923. A description of the
famous controversy (the Water Controversy) over the priority
of the discovery of the composition of water, involving Cav-
endish, Watt, Priestley, and Lavoisier.
USEFUL IDEAS DEVELOPED
1. Analysis is the breaking down of a compound into
simpler substances.
2. Synthesis is the building up of a more complex com-
pound from simpler substances.
3. The decomposition of any compound by electricity is
called electrolysis.
4. Edward Morley spent more than ten years of his life in
determining the exact ratio in which oxygen and hydrogen
unite in forming water. His work is a fine example of ac-
curacy and patience in scientific research.
5. Every pure chemical compound has a definite composi-
tion. This is the law of definite proportions. C.P. refers to a
chemically pure substance.
6. The density of any substance is the weight of a unit
volume of that substance.
7. A calory is the amount of heat necessary to raise the
temperature of one gram ol water one degree centigrade.
8. The specific heat of a substance is the number of calories
necessary to raise the temperature of one gram of that sub-
stance one degree centigrade.
9. The specific gravity of any substance is the weight of
one cubic centimeter or one milliliter of that substance com-
pared with the weight of an equal volume of water.
10. Water of crystallization is the water chemically present
in certain crystalline substances.
11. An efflorescent substance loses water of crystallization
on exposure to air; a deliquescent substance absorbs water
from the air.
72 NEW WORLD OF CHEMISTRY
USING WHAT YOU HAVE LEARNED
Group A
1. Why is water considered a compound?
2. How did Cavendish synthesize H2O?
3. What is the difference between analysis and synthesis?
4. Briefly describe how Lavoisier analyzed H2O.
5. What is electrolysis?
6. (a) Make a labeled diagram of the apparatus for the
laboratory electrolysis of H2O. Indicate the direction of the
current, the cathode, and the anode, (b) At which electrode is
the H2 given off? (c) the O2?
7. What part does sulfuric acid (H2SO4) play in the elec-
trolysis of H2O?
8. Write the word-equation for the electrolysis of H2O.
9. How would you test to find out which of the gases
present in the two outer tubes of the electrolysis apparatus
is H2?
10. (a) What is the composition of H2O by volume? (b) by
weight?
11. What are the differences between C.P., V.S.P., Tech,
and Reagent chemicals?
12. In connection with the study of H2O, cite an example of
the accuracy and patience ol men of science.
13. State and illustrate the law of definite proportions.
14. (a) Does vigorously boiling H2O have a higher tempera-
ture than slowly boiling H2O? (b) Explain.
15. What are five physical properties of H2O?
16. How does steam differ from water vapor?
17. Why is it sometimes unsafe to purchase drugs or medi-
cines by trade-mark alone?
18. Explain why water pipes often burst in cold weather.
19. (a) Which would you prefer to heat your feet on a cold
night — a hot flatiron or a bottle of hot H2O? (b) Explain.
20. If H2O were less stable than it is, what disaster might
occur?
WATER 73
21. Describe two chemical properties of H0O.
22. Give an example of the part that a trace of H2O may
play in bringing about a chemical change.
23. Illustrate what is meant by water of crystallization.
24. Transparent crystals of washing soda change to a white
powder. Is this a physical or a chemical change?
25. Why is an unstoppered bottle of calcium chloride
(CaCl2) sometimes left in large clocks?
26. When crystals of table salt (NaCl) are heated, some
H2O is liberated. Is this water of crystallization? Explain.
27. Examine some pellets of NaOH that have been exposed
to air. What property does NaOH have?
28. (a) Which gives a more severe burn, boiling H2O or
steam? (b) Why?
29. Explain the importance ol H2O to health.
Group B
30. If H2O did not expand on Iree/ing, how would it affect
us?
31. A spark is passed through a mixture ol 60 ml. of O2 and
50 ml. of H2 in the presence of water vapor. What substances,
and how much of each, will be found in the tube after the
explosion?
32. Devise an experiment to show the composition of H.,O
by weight. Mention actual weights.
33. How would you determine the percentage of H.,O
present in a sample of "dry" wood?
34. A liter of H0 weighs 0.09 g. and a liter of O2 weighs
1.43 g. Show how you could find the composition of H2O by
weight from these facts (data) .
35. What weight of oxygen can be obtained from the elec-
trolysis of 50 pounds of water?
36. Ice is purer than water. Would it be safe to use ice from
a polluted pond in your iced tea? Explain.
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Find out from your mother or grocer the cost of pound
packages of powdered *washing soda and crystallized washing
74 NEW WORLD OF CHEMISTRY
soda. The water of crystallization present in crystallized wash-
ing soda is equal to approximately 63 percent of the weight
of the crystal. On this basis, calculate which is less expensive,
the crystallized washing soda containing water of crystallization
or the powdered washing soda containing no water of crystal-
lization. Make a report of your problem in class.
2. A student came to his chemistry teacher very excited
about his invention of a system that would run an airplane
indefinitely on a small amount of H2O. He planned to de-
compose H2O by electricity, use the O2 and H2 thus produced
to supply an oxyhydrogen torch which, when ignited, would
boil H2O for a steam engine. The steam engine in operation
would give power to the plane's propeller and, at the same
time, to an electric generator. The electricity thus produced
would be harnessed to decompose more H2O, which would be
constantly renewed as a product of the burning of the O2 and
H2 in the oxyhydrogen torch. What do you think of this in-
vention? Your answer should be complete with scientific
reasons.
3. Benjamin Thompson, one of the greatest of early Amer-
ican scientists and who later became Count Rumford, made
classic contributions in the field of specific heat. "It is a strange
coincidence," wrote Albert Einstein, "that nearly all the funda-
mental work concerned with the nature of heat was done not
by professional scientists but by men who regarded science as
their great hobby." Mention two other scientific contributions
made by men who were not professional scientists.
4. Get some anhydrous copper sulfate and with it determine
whether certain "dry substances" really contain water. Report
your results.
THE BRICKS OF THE UNIVERSE
The bodies which time and nature
add to things little by little, . . .
no exertion of the eyesight can be-
hold; and so, too, wherever things
grow old by age and decay, and when
rocks hanging over the sea are eaten
away by the gnawing salt spray, you
cannot see what they lose at any
given moment. Nature, therefore,
works by unseen bodies.
Lucretius, 99-55 B.C.
The origin of the idea of atoms. The most fruitful scientific specu-
lation that came out of ancient Greece 2500 years ago was that mat-
ter is made up of small, eternal particles in continual motion. Leu-
cippus and his pupil Democritus (de mok'ri tws) , taught that all
matter was composed of invisible, indivisible, indestructible parti-
cles, or atoms. During the seventeenth century, Newton had similar
ideas about the nature of matter. "It seems probable to me," wrote
Newton, "that God in the beginning formed matter in solid, massy,
hard, impenetrable, moveable particles, so very hard as never to wear
or break into pieces." The gradual development of the idea of the
atom is an interesting story.
How Dalton's approach to the nature of matter differed from that
of the Greeks. Early in the nineteenth century, an English scientist,
John Dalton, became interested in the idea of atoms. Newton's idea
was a beautiful one, thought Dalton, but did it check with the known
facts? Would it help explain some of the physical properties of gases,
which had so puzzled him?
It is interesting to note the difference in the way the Greek teach-
ers had come to their ideas and the way in which Dalton formulated
his theory. The Greek teachers made a few observations, followed.
75
76
NEW WORLD OF CHEMISTRY
some logical reasoning, and then ventured an opinion. For example,
they noticed that a lump of salt could be broken down into bits of
salt, which could then be further reduced to tiny crystals of salt. It
was inconceivable, they reasoned, to continue this division forever.
There must come a time, therefore, when one would finally reach the
smallest piece of salt, that is, an indivisible atom of salt. Dal ton did
more. He also experimented. He tried to find out why the gases of
the atmosphere remained mixed, how gases dissolve in water, and
whether the composition of pure compounds varies or is constant.
On the basis of some not too accurate observations in these experi-
ments and some indirect evidence from facts known in his day,
he formulated the atomic theory. In 1803 he announced it suddenly
without waiting to test all of it by experiment. Since that time a
great many additional experiments have proved that Dalton's theory
was substantially correct. , ;
With this electron microscope,
scientists may study molecules
and other infinitesimal par-
ticles. Typical photomicro-
graphs may be seen in the
background.
ATOMS 77
Dalton's atomic theory. Dalton's atomic theory was based on the
following ideas:
1) That all matter consists of extremely small particles, called
atoms.
2) That all the atoms of any one element are alike in size, shape,
and weight.
3) That the atoms of one element differ from those of all other
elements in size, shape, and weight.
4) That chemical changes are changes in the combination of atoms
with each other.
5) That even in the most violent chemical change, the indivisible
atoms do not break into pieces.
The diameter of the hydrogen atom is about 1/250,000,000 inches.
This is several hundred times smaller than the average-sized bac-
terium.
To explain his theory, Dalton drew pictures of atoms. Each atom
was represented by a circle. Since the atoms of elements are unlike,
he varied the pictures of the circles as follows:
Hydrogen Oxygen Nitrogen
Carbon 0 Sulfur Q Gold
Dalton pictured chemical change as the union of one or more
atoms of one element with atoms of other elements. He believed that
when mercury is heated in air, one atom of mercury unites with one
atom of oxygen, forming a particle of the compound, mercuric oxide.
To demonstrate this union of atoms, Dalton constructed model
spheres, bringing them into contact with each other:
1 atom of liquid 1 atom of gaseous 1 molecule of red
mercury "*" oxygen * mercuric oxide
According to Dalton, atoms preserve their individuality in all
chemical changes. Hence, Dalton described an atom as the smallest
part of an element that takes part in a chemical change without itself
being altered. Atoms combine to form molecules. Two or more atoms
may combine into a molecule of an element or into a molecule of a
compound. A molecule is the smallest part of a compound or element
that has the chemical properties of that compound or element.
78 NEW WORLD OF CHEMISTRY
Inertia in scientific thinking. Dalton's theory was strongly at-
tacked by the leading scientists of his day. One of the most eminent
of them said he could not understand "how any man of sense or sci-
ence would be taken in by such a tissue of absurdities." Dalton's
theory was the result of creative imagination and the boldness of a
great thinker. Dalton had never seen nor weighed an atom. Yet his
theory was of practical value and was accepted gradually by the scien-
tific world as a useful working hypothesis by which chemical changes
could be explained.
The newer electron theory of matter (see Chapter 11) has ex-
panded Dalton's theory. It has been modified in details, but its gen-
eral applications still hold.
The use of Dalton's theory. One of the questions that scientists of
Dalton's time were studying was whether the composition of a com-
pound is always the same or whether it varies. Some believed that
compounds are always formed from fixed amounts of elements. They
believed therefore, that the composition of a compound is always the
same.
The French chemist, Claude Berthollet (ber'to'le') ran some ex-
periments to test this question. On the basis of his experiments he be-
lieved that the composition of compounds might vary to some extent.
Joseph Proust (prdost) , another Frenchman, set out to settle this
difference of opinion. He repeated Berthollet's experiments, using
the purest of chemicals and the most delicate apparatus available.
Taking every precaution to prevent error, he found mistakes in
Berthollet's work. He found that his fellow-scientist had used im-
pure compounds and substances such as glass and mixtures of metals
(alloys) and mixtures of liquids, which were not pure compounds.
For eight long years, the difference of opinion persisted. Never,
however, did it become anything but an honest, truth-seeking discus-
sion. Personal whims and prejudices did not decide the matter.
When Berthollet considered Proust's evidence and discovered his
own errors, he accepted Proust's verdict and agreed that the compo-
sition of compounds is always the same.
The law of definite proportions. This law states that the elements
in a compound always occur in a definite proportion by weight. This
is another way of saying that the composition of compounds is always
the same.
Dalton's little circles very neatly explained the law. The weight
of atoms of any element is always the same. Compounds are com-
posed of these minute and unchangeable atoms. Therefore the com-
position of compounds by weight must be definite and uniform.
ATOMS
79
Dalton discovers the law of multiple proportions. Dalton knew
that one atom of carbon • unites with one atom of oxygen O to
produce the deadly gas, carbon monoxide •O • In this compound
the carbon weighs f as much as the oxygen. This fraction can be ex-
pressed as the ratio of 3 to 4.
Carbon also combines with oxygen to form carbon dioxide. Dal-
ton wrote this combination as 0QO- In this compound the carbon
weighs only f as much as the oxygen. The ratio is three parts of car-
bon to eight parts of oxygen: 3 to 8. From observing this and other
similar combinations, Dalton formulated another fundamental law
of chemistry, the law of multiple proportions.
You note that in both carbon monoxide, CO, and in carbon di-
oxide, CO2, the weight of the carbon is the same. But the weight of
the oxygen in carbon monoxide is 4, and in carbon dioxide is 8.
Thus three parts of carbon combine with either 4 parts or 8 parts of
oxygen.
When two elements combine to form more than one compound,
with the weight of one element remaining fixed, the ratios of the
weights of the other elements are small whole numbers.
Thus the amounts of oxygen that unite with three parts of carbon
are in the ratio of 4 to 8, or 1 to 2.
How the discovery of hydrogen peroxide helped to uphold the law
of multiple proportions. In 1818, Louis Thenard (ta nar') , a French
teacher of chemistry, discovered a compound, which upon analysis
was shown to be made up of equal volumes of oxygen and hydrogen.
This compound is hydrogen peroxide.
Hydrogen and oxygen combine to form two different compounds.
Water is composed of one part hydrogen and eight parts oxygen by
weight. Hydrogen peroxide is composed of one part hydrogen and
sixteen parts oxygen. Thus the ratio of the weights of oxygen that
combine with a fixed weight of hydrogen is 8 to 16, or 1 to 2.
COMPOUND
ELEMENTS
BY WEIGHT
RATIO
( Water H2O
(^ Hydrogen peroxide H2O2
1 H
1 H
80 "I
160 /
1 to 2
f Carbon monoxide CO
^ Carbon dioxide CO2
3C
3C
40)
8O )
1 to 2
This maze of tanks and pipes is re-
quired for the commercial preparation
of hydrogen peroxide.
Properties and uses of hydrogen peroxide, FLO,. Since water and
hydrogen peroxide have different chemical compositions, they have
different physical and chemical properties. Hydrogen peroxide is a
colorless liquid, about one and one-half times as heavy as water. It
is odorless and mixes with water, alcohol, or ether. It is useful com-
mercially because it is unstable. That is, heat or light decomposes it
easily into water and oxygen:
H2O2 -* H20 + O T
The oxygen atom that is liberated is in a very active state, i^eady to
combine with another atom of oxygen or with any other substance
at the instant of liberation. This very active atomic oxygen is some-
times called nascent (newborn) . Nascent oxygen is written as (X
ordinary oxygen as O2. The arrow pointing upwards represents a gas.
Some colored compounds lose their color when they are oxidized.
Fibers containing compounds which give them their color can be
bleached by exposing them to nascent oxygen. Hydrogen peroxide
is used as an oxidizing agent to bleach, or decolorize, cotton goods,
wool, wood pulp, wood used for furniture, as well as silk, hair, feath-
ers, glue, and other animal substances.
Some bacteria are destroyed when exposed to oxygen. Hydrogen
peroxide is therefore used as a household antiseptic. The household
product actually is mainly water with a small amount (usually 3 per
cent) of hydrogen peroxide dissolved in it. It also contains some sub-
stance such as acetanilid, which retards the decomposition of the hy-
drogen peroxide. As an antiseptic, hydrogen peroxide is not very
effective since the oxygen it releases does not reach enough of the
bacteria.
80
ATOMS
81
Commercial preparation of hydrogen peroxide. It barium perox-
ide, a white solid, is treated with dilute sulfuric acid at a temperature
below Lr)°C., hydrogen peroxide and barium sultate are formed. Bar-
ium sulfate is a white insoluble solid which settles out.
Sulfuric acid -f barium peroxide — > hydrogen peroxide 4- barium sulfate
;H2!SO4~~~+"~~~BaO2j -> H2O2 + BaSO4 1
This type of reaction is called double replacement — barium replaces
hydrogen. The arrow pointing downward after BaSO4 indicates that
this compound is insoluble and separates out, or precipitates. The
insoluble substance that separates out is called a precipitate.
Most of the hydrogen peroxide produced commercially is made by
gently heating pcrsulfuric acid, H_,S,ON, which reacts with water:
Persulfuric acid
H,S>0«
+
4-
water
2H20
sulfuric acid
2H2S04
+ hydrogen peroxide
+ H2O2
In this process the hydrogen peroxide is distilled out. Superoxol is a
30% H2O2 solution. A 90% solution of this substance is used as a
rocket fuel.
The Glenn Martin Company
Off to an altitude of 160 miles roars
the Navy's Viking No. 1 1 rocket.
In flight this rocket will attain a
speed of 4300 miles per hour.
Rocket development depends
greatly upon fuel research carried
on in chemical laboratories.
82 NEW WORLD OF CHEMISTRY
YOU WILL ENJOY READING
Harrow, Benjamin. The Romance of the Atom, pp. 27-33.
Boni & Liveright, New York, 1927. Origin and development
of ideas about atoms.
Langdon-Davies, John. Inside the Atom. Harper & Bros.,
New York, 1933. Amusing, popular introduction to science
and the nature of matter.
Leicester, H. M. and Klickstein, H. S. Source Book in Chem-
istry, pp. 215-220. McGraw-Hill Book Company, New York,
1952. Gives John Dal ton's observations on the constitution of
bodies.
Thomson, J. Arthur. The Outline of Science, pp. 245-253.
G. P. Putnam's Sons, New York, 1937. "Foundations of the
Universe." In "The World of Atoms" the size and energy of
these tiny particles are very simply discussed.
USEFUL IDEAS DEVELOPED
1. The idea of individual indivisible particles of matter
originated with the ancient Greek teachers. Dalton used this
[idea and developed the atomic theory in 1803.
2. The chief assumptions of Dalton's theory were: a) That
all matter consists of extremely small particles, called atoms,
b) That all the atoms of any one element are alike in size,
shape, and weight, c) That the atoms of one element differ
from those of all other elements in size, shape, and weight,
d) That chemical changes are changes in the combination of
atoms with each other, e) That even in the most violent
chemical changes, the indivisible atoms do not break into
pieces.
3. An atom is the smallest part of an element that takes
part in a chemical change without itself being altered.
4. A molecule is composed of two or more atoms and is the
smallest part of a compound or element that has the proper-
ties of that compound or element.
5. Personal whims, prejudices, or prestige should play no
part in settling scientific problems.
6. The law of definite proportions states that the elements
in a compound always occur in a definite proportion by
weight.
7. The law of multiple proportions states: When any two
elements combine to form more than one compound, with the
weight of one element remaining fixed, the ratios of the weights
of the other element are small whole numbers.
ATOMS
83
USING WHAT YOU HAVE LEARNED
Group A
1. When did John Dalton advance his atomic theory?
2. Was the idea of atoms new in Dalton's time? Explain.
3. How did Dalton's approach to the study of the nature of
matter differ from that of the Greeks?
4. State the five essentials of Dalton's atomic theory.
5. How did Dalton distinguish between atoms of different
elements?
6. According to Dalton, what happens when elements
unite?
7. How did the scientific world receive Dalton's theory?
8. Describe the difference of opinion between Berthollet
and Proust.
9. (a) State the fundamental law of chemistry that Dalton
discovered.
(b) Show how it applies to the composition of CO and
C00.
10. (a) When and (b) by whom was H2O2 prepared for the
first time?
1 1 . Compare the physical properties of H2O and H2O2.
12. Contrast the chemical properties of H2O and H2O2.
13. Compare the compositions of H2O and H2O2 by volume
and weight.
14. State two ways in which the decomposition of H2O2 may
be retarded.
15. Write the word-equation for the decomposition of H2O2.
16. What is the difference between O2 and nascent oxygen?
t . . .
17. Describe the two most useful properties of H2O2.
18. What is the great advantage of H2O2 as a bleaching
agent?
19. What is an antiseptic?
20. State a commercial method of preparing H2O2.
21. Write the equation for the preparation of H2O2 from
BaO0.
84 NEW WORLD OF CHEMISTRY
22. What is a precipitate?
23. Compare the chemical properties of O2, O3, and O.
Group B
24. Explain how the law of multiple proportions is based
on Dalton's atomic theory.
25. Two oxides of nitrogen, nitrous oxide (N.,O) and ni-
trogen dioxide (NO2) , show the following ratios by weight:
In nitrous oxide, the ratio by weight of nitrogen to oxygen
is 7 to 4; in nitrogen dioxide, the ratio by weight of nitrogen
to oxygen is 14 to 32. With what fundamental law of chemistry
are these figures in accord?
26. Do we still believe that all atoms of the same element
weigh the same?
27. (a) Is Dalton's atomic theory still a theory or has it been
proved experimentally? (b) Explain your answer.
28. Dalton's hobby of recording weather conditions was
greatly responsible for the atomic theory. Can you cite another
example in science where a hobby has resulted in a great
advance?
29. Write a 2- or 3-page report on the life of John Dalton.
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Dalton's atomic theory is a beautiful example of creative
imagination in pure science. Until very recently America, al-
though it excelled Europe in inventions and applied science,
has lagged behind the Old World in the kind of creative imagi-
nation represented by Dalton's theory and the germ theory of
disease which Pasteur gave humanity. Can you give reasons
for this state of affairs? What of the future?
2. Add a pinch of manganese dioxide to one-third of a test
tube of hydrogen peroxide from your medicine cabinet. Insert
a glowing splint. Prepare a brief report on the reaction which
takes place.
3. Prepare a debate or a class discussion on the topic, "If
prejudices and superstitions did not exist, social change would
proceed more rapidly."
4. Form a committee to investigate and report on prejudices
and superstitions common to your community. Do any of these
have a scientific foundation?
5. Make models of the atoms of hydrogen, carbon, sulfur,
gold, mercury, and nitrogen as Dalton constructed them, using
marbles or some other suitable spheres.
FORMULAS:
THE CHEMIST'S
ABBREVIATIONS
The mysterious symbols of alchemy. Ancient alchemists feared
that they would lose their unique social position if anyone could
understand their work. Consequently they used strange symbols,
both to conceal the true nature of their writings and also to give
themselves an added air of mystery and magic. They represented
sand by .jjj.^ , glass by Q— Q, soap by ^^-, and salt by ©. To
• • • • • ^^^^^ ^"^
them the symbcJ for perfection and also for the sun was the circle
O • Hence it was used to represent sun-colored gold, which they con-
sidered the perfect metal.
Many of their symbols were derived from ancient mythology. The
lance and shield of Mars, god of war, £f , represented, appropriately
enough, iron. The looking glass of Venus, O, was their symbol for
copper because, according to legend, Venus had first appeared on the
shores of the island of Cyprus, long famous for its copper mines.
Often an alchemist would develop a set of symbols for his own
special use and would reveal the meaning to no one, not even to his
brother alchemists. In one Italian manuscript written in the seven-
teenth century, the element mercury was represented by 20 different
symbols and 35 different names. As long as alchemy was a purely
personal practice carried on for selfish ends, the confusing symbolism
presented no problem. However, when the universal science of
modern chemistry began to emerge, it was essential to develop a
system of symbols which could be understood easily in every country
of the world. Dalton, and a few others before him, had attempted
to substitute some reasonable system for the jungle of weird signs
and strange names used by the alchemists. However, these attempts
failed — largely because they were unwieldy and inconvenient.
85
86 NEW WORLD OF CHEMISTRY
Berzelius helps bring order out of chaos. Jons Berzelius (ber-
ze'li-ws) was a Swedish orphan who became, in the words of Sir
Humphry Davy, "one of the great ornaments of his age." Berze-
lius invented a simple system of chemical notation which he intro-
duced in 1814. Today it is used in every country of the world.
Berzelius said very logically: "It is easier to write an abbreviated
word than to draw a figure. The chemical signs ought to be letters
for the greater ease of writing and not to disfigure a printed book.
I shall therefore take for the chemical sign the initial letter of the
Latin name of each chemical element, thus, C, H, N, O, S, and P.
If the first letter be common to two metals, I shall use both the initial
letter and another letter they have not in common, as gold (aurum) ,
Au; silver (argentum) , Ag; antimony (stibium) , Sb; tin (stannum) ,
Sn." The first letter of the symbol is always capitalized; the second,
if there is a second, is not.
How a compound is represented by a formula. When we are deal-
ing with a compound rather than an element, we write its abbrevi-
ated form, or formula, by placing side by side the symbols of the
elements that compose the compound. For example, the formula of
zinc oxide, a compound of zinc and oxygen, is ZnO; the formula of
hydrochloric acid, a compound of hydrogen and chlorine, is HC1.
A formula not only represents the name of a compound, but also
one molecule of that compound. Similarly a symbol represents both
the name of an element, and one atom of that element.
The use of the subscript. Berzelius used a subscript, a small num-
ber placed below and to the right of a symbol, to indicate the number
of atoms of the element represented by that symbol. Thus, H2O
represents one molecule of water containing two atoms of hydrogen
and one atom of oxygen. The subscript 1 is never written.
As you use Berzelius' system, you will see its remarkable simplicity
and great value. However, it was not accepted without a struggle.
Even Dalton protested against it saying, "Berzelius' symbols are
horrifying. A young student might as soon learn Hebrew as make
himself acquainted with them." Evidently, Dalton must have for-
gotten his own very complex pictures of various compounds.
How is the proper subscript determined? You may have already
wondered how it is possible to tell what subscript to use. Why A12O8
and ZnCl2 and not A1O and ZnCl? Must we memorize all formulas or
are there definite rules to guide us?
Fortunately, it is possible to write formulas without first mem-
orizing them. To do so, however, we must have a thorough under-
standing of valence.
FORMULAS
87
What is valence? The law of definite proportions tells us that
atoms occur in compounds in fixed ratios. The following table shows
the ratio in which hydrogen combines with other elements to form
four common compounds.
COMPOUND
NUMBER OF
HYDROGEN ATOMS
IN COMPOUND
NUMBER OF ATOMS
OF SECOND ELEMENT
IN COMPOUND
HCI (hydrogen chloride)
H2O (water)
NH3 (ammonia)
CH* (methane)
1
2
3
4
1 chlorine atom
1 oxygen atom
1 nitrogen atom
1 carbon atom
In this table, note the difference in the number of hydrogen atoms
with which one atom of the other elements combines. The valence
of an element is the number that tells us how many atoms of hydro-
gen normally combine with one atom of that element. This is a
simplified definition, but will serve us for the present.
From the table we see that one atom of chlorine combines with
one atom of hydrogen to form HCI. Thus, we say that chlorine has a
valence of one, or as chemists put it, is monovalent. One atom of
oxygen combines with two atoms of hydrogen. So we say that oxygen
has a valence of two, or is divalent. Nitrogen has a valence of three,
or is trivalent. Carbon has a valence of four, or is tetravalent.
The idea of valence was introduced in 1852 by Edward Frank-
land, an English chemist. Hydrogen is used as the standard, because
its atom never combines with more than one atom of any other ele-
ment. Hence, if its valence is considered to be one, the valence of
every other element must be a whole number.
Learning valences. There is no "royal road" to the study of va-
lence. You will find that memorizing the valences of the more com-
mon elements will save you a tremendous amount of work and will
give you a better understanding of the material to follow. Table 2
lists the valences which are considered essential.
Notice th'at in the table, the symbols of the metals are followed
by plus (+) signs and the nonmetals and radicals followed by
minus signs (•— ) . These signs represent electric charges, for reasons
which you will learn in Chapter 16. The number of these electric
charges correspond to the valences of the element or radical. Group-
88
NEW WORLD OF CHEMISTRY
ing elements as metals and nonmetals is not a completely satisfactory
means of classification because some elements behave at times as
metals and at other times as nonmetals.
Binary compounds. A binary compound is composed either of two
elements, two radicals, or one element and one radical. If you keep
in mind the following rules, you will find writing the formulas for
binary compounds is a simple process. Practice is essential.
IMPORTANT VALENCES
K;'
f
MONOVALENT
DIVALENT
TRIVALENT
1
i
I METALS
*Ammonium NH4 *
Copper (cuprous) Cu +
Lithium Li *
Mercury (ous) Hg *
Potassium K
Silver Ag +
Sodium Na +
Barium Ba "*"*"
Calcium Ca++
Copper (cupric) Cu **
Iron (ferrous) Fe **
Magnesium Mg"1"1"
Mercury (ic) Hg*^
Zinc In**
Aluminum Al *++
Antimony Sb"1"*"*1
Arsenic As * <f+
Chromium Cr "*"*"*'
Iron (ferric) Fe "*"*"*"
NONMETALS
Bromine (bromide) Br~
Chlorine (chloride) Cl~
Fluorine (fluoride) F "
Iodine (iodide) l~
Oxygen (oxide) O
Sulfur (sulfide) S~"
Nitrogen N
(nitride)
Phosphorus P
(phosphide)
RADICALS
*Acetate C2H3O2~
Bicarbonate HCOs"
Chlorate CI03"
Hydroxide OH~
Nitrate NOa "
Nitrite NO2 -
Carbonate COs
Sulfate S04~~
Sulfite SO 3 —
Phosphate PO4
* Radicals: A radical is a group of atoms acting as a single atom and having its awn individual
valence. The ammonium radical, NH^ has a valence of ?.
!,
HOW TO USE VALENCE IN WRITING FORMULAS OF
BINARY COMPOUNDS
1 ) Write the symbol with a positive valence first, followed by
the symbol with a negative valence. Add plus and minus signs
to the upper* right of the symbols, that is, in superscripts, to
show the valence for each symbol.
2) // the valences of the symbols are equal no subscripts are
added. This rule is followed unless the subscripts represent
the actual structure of the compound. Thus, the formula of
hydrogen peroxide is H2O2 and not HO, since a molecule of
hydrogen peroxide actually contains two atoms of hydrogen
and two atoms of oxygen.
FORMULAS 89
3) Since every compound is electrically neutral, the number
of positive charges must be the same as the number of its
negative charges. Therefore, // the valences are not equal in
numerical value, subscripts must be added to equalize them.
Add to each symbol a subscript of the same numerical value
as the valence of the other symbol. The subscript 1 is never
written.
4) A radical acts like an element, that is, it usually passes
thiough a chemical reaction unchanged. It should be placed
in parentheses only if it is followed by a subscript greater than
one.
EXAMPLE A: Write the formula for the compound, zinc oxide.
1) Zinc is written Zn and has a valence ot plus two; oxygen
is written O and has a valence of minus two. The symbol lor
zinc appears first in the formula since zinc is a metal and oxy-
gen is a nonmetal. Indicate valences by plus and minus signs.
Zn++O—
2) Since the valences are equal no subscripts are written,
and the subscript for each symbol is understood to be one.
3) The proper formula for zinc oxide therefore is ZnO.
EXAMPLE B: Write the formula for the compound, cupric
chloride.
1) Copper is written Cu and has a positive valence of two;
chlorine is written Cl and has a negative valence of one. Write
the symbol for copper first because it has the positive valence.
Indicate the valence of each element by using plus or minus
signs. (Note that the cupric valence of copper should be used.
See Table 2.)
Cu++Cl-
2) The subscript of each symbol must be equal to the va-
lence of the other symbol. Since Cu has a valence of plus two,
give Cl the subscript two. Since Cl has a valence of minus one,
the subscript for Cu is understood to be one. The crossed ar-
rows show these relationships.
The negative and positive valences are equal because there is
one atom of copper with a valence of plus two, and two atoms
of chlorine each with a valence of minus one.
3) The proper formula for cupric chloride is therefore
CuCU.
90 NEW WORLD OF CHEMISTRY
EXAMPLE C: Write the formula for the compound, magnesium
sulfate.
1 ) Magnesium (Mg) has a valence of plus two. The sulfate
radical (SO4) has a valence of minus two. Therefore, Mg ap-
pears first in the formula. Indicate the valences by using plus
and minus signs.
Mg++SO4—
2) The valences of the two symbols are equal. Therefore,
the subscript of each is understood to be one, and no sub-
scripts are written.
3) The proper formula for magnesium sulfate is MgSO4.
EXAMPLE D: Write the formula for the compound, zinc nitrate.
I ) Zinc (Zn) has a valence of plus two. The nitrate radical
(NO3) has a valence of minus one. Therefore, Zn appears first
in the formula. Indicate the valences by using plus and minus
signs.
Zn++NO3-
2) Since Zn has a valence of plus two, we give NO3 the sub-
script two and enclose it in parentheses. NO3 has a valence of
minus one and we consider Zn to have the subscript one. The
crossed arrows show these relationships.
3) The proper formula for zinc nitrate is Zn(NO3)2.
Some elements have more than one valence. Iron is divalent
in ferrous compounds and trivalent in ferric compounds. Fer-
rous chloride provides an example of divalent iron: Fe++Cl2-.
Ferric chloride provides an example of trivalent iron:
Fe+++Cl3-.
Mercury is monovalent in mercurous compounds and di-
valent in mercuric compounds; copper is monovalent in
cuprous compounds and divalent in cupric compounds. What
is the meaning of -ous and -ic as related to valence?
In the formula Fe3O4 (Magnetic oxide of iron) you might
think that iron has a valence of four and oxygen a valence of
three. However, the real explanation is that this compound
is a combination of Fe++O~-, in which Fe is divalent, and
Fea+++O8- ~, in which Fe is trivalent. Oxygen is always divalent.
Or, Fe364 may be thought of as FeO • Fe2O3. From these
formulas it is apparent that some elements may exhibit two or
more valences. The electron theory offers an interesting
FORMULAS 91
explanation of the fact that certain elements have more than
one valence (see Chapter 11).
PRACTICE IN WRITING FORMULAS
1. Write the formulas for the following compounds showing
the + and — signs and the arrows pointing toward the sub-
scripts: (a) sodium chloride, (b) calcium bromide, (c) ferric
iodide, (d) potassium fluoride, (e) barium oxide.
2. Write the formulas of: (a) magnesium chloride, (b) zinc
oxide, (c) aluminum nitride, (d) potassium sulfide, (e) cu-
prous chloride.
3. Write the formulas of: (a) aluminum chloride, (b) arse-
nic oxide, (c) calcium phosphide.
4. Write the formulas of: (a) sodium hydroxide, (b) potas-
sium sulfate, (c) mercurous phosphate.
5. Write the formulas of: (a) calcium bicarbonate, (b) cu-
prous carbonate, (c) magnesium phosphate.
6. Write the formulas of: (a) aluminum nitrate, (b) ferric
sulfate, (c) chromium phosphate.
7. Carbon (C++++) and silicon (Si++++) are tetravalent.
Write the formulas of: (a) carbon tetrachloride and (b) sili-
con dioxide.
8. Write the formulas of: (a) mercuric nitrate, (b) sodium
nitrate, (c) mercuric chlorate, (d) mercuric hydroxide, (e)
mercuric carbonate, (f) mercurous sulfate, (g) calcium sul-
fite, (h) mercuric phosphate, (i) mercurous chloride.
How to determine valence in compounds of more than two ele-
ments. Remembering that in every compound the number of posi-
tive charges must equal the number of negative charges, let us find
the valence of Cr in K2CrO4. There are four oxygen atoms each with
two negative charges making a total of eight negative charges. To bal-
ance these, the compound has two potassium atoms each having one
positive charge or a total of two positive charges. The compound
must therefore have six more positive charges which must come from
the metal chromium. Hence the valence of Cr in this compound is
plus six.
Chemistry has a language and nomenclature of its own. Lavoisier
realized the importance of language to a science. In 1789, the year
in which the Bastille was stormed, he published a book written in
the new language of chemistry. It did not contain the obscure words,
the mystic symbols, and the pompous phrases of alchemy.
An outstanding nuclear scientist is Dr. Glenn T.
Seaborg of the Radiation Laboratory, University of
California, Berkeley, California. He has played an
important role in the discovery of the trans-uranium
elements, numbers 93-100. What names have
been given elements 93-98? Can you suggest
their derivation?
University <>! Call In
In naming the elements, several methods were used. Some, in-
cluding bromine, meaning stench, were named after a physical prop-
erty. Some, including argon, meaning idle, were named after a chem-
ical property. Some, including /;o/onium, germanium, gallium, and
americium, were named after countries or other geographic regions.
Some were named after the city or state connected with their dis-
covery. Thus hafnium was christened after the Latin name for the
city of Copenhagen, where it was discovered. Radioactive curium was
named after the Curies, who were the earliest investigators in the field
of radioactivity. 77«0rium and tantalum were named after figures
in mythology. The origins of the names of the earliest known ele-
ments have been lost in the darkness of antiquity.
Names of metals and metallic radicals usually end in either -turn
or -um, as sodium, potassium, platinum, curium, hafnium, alumi-
num, calcium, and the ammonium radical.
Names of nonmetals and nonmetallic radicals usually end in -ine
or -gen, as chlorine, bromine, iodine, oxygen, and nitrogen.
The most common suffixes used in naming compounds are -ide,
-ate, -ite, -ous, and -ic. The suflix -ide represents a binary compound.
The suffixes -ate and -ite indicate, as a rule, compounds of three
elements, one of which is oxygen. The -ite compound contains fewer
oxygen atoms than the corresponding -ate compound. Thus sodium
sulfate, Na.,SO4, contains four atoms of oxygen, and sodium sulfite,
Na.,SO3, three atoms of oxygen. The -ate compounds are salts (see
page 197) of -ic acids, and -ite compounds are salts of -ous acids.
The suffixes -ous and -ic indicate compounds in which the metal
has a lower valence in the case of -ous and a higher valence in the case
of -ic. An -ous acid contains fewer oxygen atoms than an -ic acid;
thus sulfurous acid is H2SO3 and sulfuric acid is H2SO4; chlorous
acid is HC1O., and chloric acid is HC1O3.
92
FORMULAS 93
Some commonly used prefixes in chemistry are mono- (or uni-) ,
di- (or hi-) , tri-, tetra-, and pent-. Mono- (or uni-) , di- (or hi-) , tri-,
tetra-, and pent- stand for one, two, three, four, and five atoms. Thus,
CO is carbon monoxide, and CO2 is carbon dioxide. P2O3 is phos-
phorus trioxide and P2O5 is phosphorus pentoxide. The prefix per-
means more atoms of an element than are found in a more common
compound, and the prefix hypo- means less atoms. Thus chloric acid
is HC1O3, perchloric acid is HC1O4; chlorous acid is HC1O2, and
hypochlorous acid is HC1O.
In organic chemistry, a branch of chemistry which deals with the
more than 650,000 compounds of carbon, a more comprehensive
nomenclature has been carefully worked out. We shall learn more
about organic chemistry later.
YOU WILL ENJOY READING
Caven, Robert M. and Cranston, John A. Symbols and
Formulae in Chemistry, pp. 1-29. Blackie &: Son, London,
1928. Development and use of symbols and formulas are traced
with great clearness and interest in this valuable work.
Jaffe, Bernard. Crucibles: The Story of Chemistry, pp. 136-
156. Simon and Schuster, New Ybrk, 1948. "A Swede Tears Up
a Picture Book" deals with Berzelius' life and his contributions
to the development of chemistry.
Kendall, James. At Home among the Atoms. D. Appleton-
Century Co., New York, 1932. This eminent chemist intro-
duces in novel form the problem of valences. He calls the
chapter "Valencia."
Oesper, Ralph E. "The Birth of Modern Chemical Nomen-
clature." Journal of Chemical Education, June, 1945, pp. 290-
292. Aii interesting story.
USEFUL IDEAS DEVELOPED
1. Each element has a symbol that stands for one atom of
the element.
2. Each chemical compound may be written in abbreviated
form as a formula by placing the symbols of the elements that
compose the compound side by side.
3. A radical is a group of elements that act as a single ele-
ment.
4. Subscripts ate used in a chemical formula to indicate the
number of atoms of any element which occur in a molecule of
94 NEW WORLD OF CHEMISTRY
any substance. Subscripts are also used following the paren-
theses around radical groups to indicate how many radicals
are present in one molecule.
5. The valence of an element is a number that represents
the number of atoms of hydrogen with which one atom of that
element normally combines in forming a compound.
6. Names of metals and metallic radicals usually end in -turn
or -urn; names of nonmetals often end in -ine or -gen.
USING WHAT YOU HAVE LEARNED
Group A
1. What is a symbol?
2. Who introduced the modern symbols and formulas of
chemistry?
3. Why was it necessary to replace the old alchemical
symbols with a new system of symbols?
4. Write the symbols of three gaseous, one liquid, and
three solid elements.
5. Several elements have the same initial letter. How do
we indicate these elements by symbols?
6. What does a formula indicate?
7. (a) Define valence, (b) Who introduced valence into
chemistry?
8. What element having a valence of one is used in de-
termining the valence of other elements?
9. What kind of element would have a valence of zero?
10. Of what importance to you is a knowledge of valence?
1 1. What is a radical?
12. Copy and complete the following statements: Metals
have . . . valences; nonmetals have . . . valences. The most
common radical with a plus valence is ....
13. Copy and complete the following statements: Three
metals with a valence of one are . . . , . . . , and .... Three
divalent metals are . . ., . . ., and .... Three monovalent non-
metals are . . . , . . . , and .... The bicarbonate radical has a
valence of .... A radical with a valence of three is ....
14. State four rules for writing formulas. Illustrate each.
FORMULAS
95
15. Give two examples of elements that have more than
one valence.
16. Copy and complete the following table. Do not write
in this book.
Bromide Sulfide Chlorate Sulfate Phosphate Oxide Hydroxide
Silver
AgBr
Chromium
Mercuric
Ammonium
1
17. Make a list of the formulas of the phosphates of eight
different metals.
18. Make a list of the formulas of the carbonates of eight
different metals.
19. Make a list of the formulas of five ammonium com-
pounds.
20. Correct the following: FeCl, CuS2, Ag(NO3)2, KSO4,
Na2C103, and NH4 (OH) .
21. Write the names and formulas of five compounds of
zinc.
It ...
22. Name five elements and tell how their names were de-
rived.
23. How can you tell whether the element ruthenium is a
metal or a nonmetal?
24. Give the formulas and names of two acids that illustrate
the difference in the use of the suffixes ~ous and -ic.
25. Give the formulas and names of two compounds other
than acids that illustrate the difference in the use of -ous
and -ic.
I . .
26. How would you name the two compounds, BaO and
Ba02?
27. Explain the meaning of each letter and subscript in
these formulas: HNO8, FePO4, Na2CO8.
28. Mark the valences of the elements and radicals in these
compounds, using -{- and — signs. CuSO4, HgCl2, NaClO3,
Ca (HC03) 2, X3 (P04) 2, (NH4) 8Y, MnO2.
46 NEW WORLD OF CHEMISTRY
29. Determine the valence of: (a) sulfur in H2SO4, (b)
manganese in KMnO4, and (c) chromium in K2Cr2O7.
Group B
30. An unknown element X has a valence of three. Write
the formula for its oxide.
31. The metal Y has a valence of two and the nonmetal Z
has a valence of two. Write the formula for their compound.
32. The nomenclature of chemistry is still not completely
organized. Can you give any reasons for this state of affairs?
33. What element was named after a Finnish chemist?
(Read the chapter on the rare-earth elements in Weeks' Dis-
covery of the Elements.)
34. Write a two- or three-page report on the life of Berzelius.
See list of additional reading material.
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Make a chart and, with the help of your teacher of art,
include as many alchemical symbols as you can and their
modern equivalents.
2. The old alchemists wanted to keep their knowledge hid-
den from the rest of the world, and so used strange symbols
and mysterious language, (a) Name two groups of people who
are modern equivalents of the alchemists, (b) Tell how they
keep their knowledge to themselves and clothe their activities
with a veil of mystery, (c) What are the reasons for their
secrecy?
3. Form a committee to make a report on the trades, pro-
fessions, and businesses in your community that have "lan-
guages of their own." (a) What reasons have they for these
languages? (b) What, after all, is the function of language?
/. ATMOSPHERE:
THE OCEAN OF AIR
. . . It ought to be esteemed much
less disgraceful to quit an error for
a truth than to be guilty of the van-
ity and perverseness of believing a
thing still, because we once believed
it. Robert Boyle, 1627-1691
Exploring the atmosphere. Every day many of us see gleaming air-
planes flying swiftly overhead. What are conditions up in the air
where they travel? Of what is the air in which they travel composed?
Have you ever seen the glowing trail of a meteor as it flashed
across the sky? Why did it burn? Have you ever examined the fused,
fire-scarred surface of a meteorite? Was it hot when it entered our
atmosphere, or did it become heated as it traveled through it?
Probably you know the answers to these questions, but primitive
men would not have been able to answer them. The most learned of
the alchemists, even as late as the seventeenth century, did not have
the answers. Some of them believed that air was empty space without
weight and without substance, while others believed air to be one of
the "four elements."
It was not until 1643 that Torricelli (tor-re-chel'le) invented the
mercury barometer, which shows that air has weight and exerts pres-
sure. We know today that approximately 15 pounds (14.7 Ib.) of
air rest upon every square inch of the surface of the earth. Upon
every square foot rests a column of air weighing more than one ton!
The density of the atmosphere is not the same throughout its
entire extent. It is densest at the surface of the earth, and the farther
97
the surface of the earth, the lighter the air becomes, At five
above the earth, air pressure Is only 5.5 pounds. At this eleva-
tion the engine of an airplane cannot get sufficient oxygen to burn its
fuel. U is half-starved. What it needs to function properly is more
oxygen, The turbosupercharger supplies this by foreleg more air
into the It has been calculated that even at a height of 2000
miles there 'is still some air, although above five miles normal breath-
ing of human is impossible. About 90 percent of the weight
of°thf is within 12 miles of the earth,
With respect to temperature, the atmosphere Is divided into sev-
eral layers, 'The lowest layer, about six miles deep, is called the
troposphere. Within the troposphere the temperature falls about
: 1°F. with each, rise of 300 feel. The layer above this is called the.
! stratosphere where temperature from about —30° to — 65°F.
I In the next layer, which lies between 18 and 28 miles above the
earth's surface, a temperature of about 65 degrees below zero oi^the
Fahrenheit scale is almost constant. Above these layers, at a height
of between 50 and 200 miles, are several layers of electrified particles
(of which the Kennelly-Hcaviside and the Appleton were the earliest
discovered) , which prevent most radio waves from out into
space, and" reflect them back to the earth on their way around the
globe. This is called the ionosphere, Its from.
A of a cross-section of the showing fhit approximate relative"
osition of the various layers.
ATMOSPHERE 99
morning to night and varies with the seasons. The temperature in
these ionized bands rises with elevation.
Extremely short radio waves of sufficient energy can penetrate
these electrified layers, as was proved in 1946 by Army Signal Corps
scientists, who succeeded in sending radar impulses to the moon.
Reflections from the moon were received. The round trip, 477,714
miles, was made in 2.4 seconds.
What is air? Air is made up of a number of different gases. During
the eighteenth century three of these — nitrogen, oxygen, and car-
bon dioxide — were isolated, or obtained separately, from it. Indi-
cations were noted that still other gases, less easily obtained in a
pure state, might possibly be present in small quantities. These
gases were actually found later on.
Is air a mixture or a compound? How are the gases of the air held
together? Are they chemically united or are they merely mixed, just
as sand might be mixed with clay? To answer these questions, sci-
entists made use of the law of definite proportions. If air is a com-
pound, they reasoned, then the composition of air must be constant.
To determine whether or not the composition of air is constant,
Dalton and others collected and analyzed samples of air taken from
thousands of different places — from the tops of mountains, over
lakes, in valleys, in sparsely settled regions, and in congested areas.
Gay-Lussac (ga-lu-sak') ascended over Paris in a hydrogen-filled
balloon to a height of 4 miles to get samples of air. Much more re-
cently, rockets carrying self-sealing bottles brought down air from
heights as much as 36 miles.
Although these analyses showed that all the samples of air varied
only slightly in composition, enough difference was noted to indicate
that the composition of air is not constant. Hence, air could not be
a compound. Air, they decided, is a mixture of gases.
Some other proofs that air is a mixture. The conclusion reached
by the early investigators that air is a mixture was strengthened by
further evidence. For example, rain water contains air that has been
dissolved from the atmosphere. By boiling rain water, this air can be
driven out and collected. Analysis shows that this air contains almost
twice as much oxygen as common air.
If air were a compound, dissolving it in water would not cause any
change in its composition. But because air is a mixture, each sub-
stance of which it is composed dissolves in water in proportion to its
own solubility therein. Since oxygen is more soluble in water than
nitrogen, the other main component of air, dissolved air contains
more oxygen than does common air.
100
white
phosphorus
eudiometer
9lass , -,
cylinder ',.
rise of water
Fig. 18. Finding the per-
centage of oxygen in the
air by the use of a eu-
diometer.
iron filings
eudiometer
"','V:'^*?V ,
rise of water
When scientists succeeded in changing air into a liquid by cooling
it to a temperature of about — 190°C., they found that liquid air
does not have a definite boiling point. Instead of all the liquid air
boiling at a definite temperature, as does water or any other pure
compound, they found that nitrogen boiled off first. Other com-
ponents of the air boiled off at higher temperatures. A pure com-
pound has a definite boiling point, but liquid air does not. This is
further proof that air is a mixture.
Finding the percentage of oxygen in air. The percentage of oxy-
gen in air is found by using a chemical which will react with oxygen
and remove it from the air. For example, we might place a small
piece of white phosphorus on the coiled end of a copper wire and
insert it in a eudiometer, or measuring tube, which is then inverted
in a cylinder of water (see Fig. 18) .
The phosphorus, which is easily oxidized, soon combines with the
oxygen of the air in the eudiometer according to the equation:
This reaction continues until the last trace of oxygen has combined,
forming phosphorus trioxide, white feathery crystals which readily
dissolve in the water in the cylinder. Phosphorus pentoxide, P2O5, is
formed also according to the equation:
This is a white solid which readily combines with water and, hence,
is used as a drying agent, especially for gases.
As the oxygen unites with the phosphorus, a partial vacuum is
formed in the eudiometer, and the greater air pressure outside the
* Also written, respectively, P4OC and P4O10.
ATMOSPHERE
101
tube forces water up the tube. At the end of a few hours the chemical
action ceases and the water stops rising. The volume of water that has
risen in the eudiometer is equal to the volume of oxygen that was
originally present in the air in the measuring tube. Because this rise
in the height of the water is approximately one-fifth the height of the
tube, this rise indicates that about 20 percent of air is pure oxygen.
In performing this experiment, do not touch the phosphorus.
Instead of phosphorus, a less active element such as iron may be
used. To perform this experiment, the inside of a eudiometer is
moistened with water and enough iron filings are added to form
a thin layer of iron on the inside walls of the tube, as shown in
Fig. 18. The tube is then inverted in water, and left undisturbed
overnight. Upon examination, it is found that most of the iron has
changed to brown rust, an oxide of iron, and the water has risen
about one-fifth the height of the eudiometer.
As some molten metals
harden in the mold, tiny
holes are formed due to
the presence of dissolved
gases. These holes make
the casting unfit for use.
To prevent formation of
holes in the casting,
nitrogen is bubbled through
the molten metal in the
mold. The nitrogen carries
away the undesirable
gases. This process is
known as n/frogen degas-
sing.
Carbide and Carbon Chemicals Company
102 NEW WORLD OF CHEMISTRY
Nitrogen is the most abundant element in the atmosphere. Nitro-
gen is the principal gas that remains after oxygen has been removed
from air. It was first carefully studied in 1772, by Daniel Rutherford,
uncle of Sir Walter Scott. Pure nitrogen may be prepared in the
laboratory by gently heating ammonium nitrite, NH4NO2.
NH4NO2 -» N2 T + 2H2O
The physical properties of nitrogen resemble those of its partner,
oxygen, quite closely. It is colorless, odorless, and scarcely soluble
in water (about two liters of nitrogen dissolve in 100 liters of water
under standard conditions) . It is even more difficult to liquefy than
oxygen, requiring a temperature 13°C. lower. It is slightly lighter
than oxygen. Dry air is about 78 percent nitrogen by volume.
Chemical properties of nitrogen. Since nitrogen under normal
conditions is not chemically active, Lavoisier named this element
azote, meaning devoid of life. However, because it is one of the ele-
ments found in niter (potassium nitrate) , this name was later
changed to nitrogen.
Chemically, nitrogen differs completely from oxygen. It does not
burn under normal conditions, and does not support respiration.
It unites with oxygen only at such temperatures as exist in the electric
arc. However, at higher and carefully controlled temperatures and
pressures, nitrogen combines with oxygen, hydrogen, and carbon.
With calcium, magnesium, lithium, and wolfram, it forms a group
of compounds called nitrides.
N2 + O2 -» 2NO (nitric oxide)
N2 + 3H2 — » 2NH3 (ammonia)
N2 + 2C — > C2N2 (cyanogen — a colorless, poisonous gas)
N2 4- 3Mg — •> Mg3N2 (magnesium nitride)
Life as we know it consists of plant and animal forms that have
survived because they were well suited to living in the earth's at-
mosphere. Life on earth might be radically different if the oxygen of
the air were not diluted with inactive nitrogen. Nitrogen tones down
the chemical activity of the oxygen of the air. Hence burning and
other oxidations are not as rapid as they otherwise would be.
How nitrogen is used. The vast storehouse of free nitrogen in air
furnishes an almost limitless supply of this valuable element. At-
mospheric nitrogen obtained from liquid air is used in making am-
monia, ammonium compounds, and nitrates. These nitrogen com-
pounds are essential in the manufacture of explosives and in
fertilizers (see pages 268, 270, 468, and 470). Nitrogen thus plays
a dual role, aiding both in supporting life and in destroying it.
ATMOSPHERE 103
Rayleigh is confronted with another riddle of the atmosphere. In
1894, Rayleigh (ra'li) , an English scientist, found that "pure" nitro-
gen obtained from air weighed a little more than an equal volume of
nitrogen obtained from pure nitrogen compounds. This puzzled him.
Dalton had declared that the weight of the atom never changed,
regardless of its source.
The difference in weight that Rayleigh noticed was very small, and
might have been ignored as caused by experimental errors. But after
Rayleigh had spent months investigating this problem, he became
convinced that the presence of some other element in the air was
responsible for the difference in the weights of the samples of nitro-
gen. With the persistence of a true research scientist he finally tracked
down the cause of this difference.
The discovery of argon, an inert gas of the atmosphere. Rayleigh,
together with William Ramsay, studied the experiments of Cav-
endish and came across the following statement: "I made an experi-
ment to determine whether the whole of the nitrogen of the at-
mosphere could be changed to nitric acid. Having condensed as
much as I could, only a small bubble of air remained. So that if
there is any part of the nitrogen of our atmosphere which differs
from the rest, we may safely conclude that it is not more than Y^th
part of the whole."
Here was a clue to their problem. Small as this quantity was,
Cavendish had not treated it as negligible or as an error in his work.
Rayleigh and Ramsay therefore repeated his experiments and isolated
a small volume of this gas from the nitrogen of the air. After sub-
jecting it to every known test, they finally identified a new element,
heavier than nitrogen, which, because of its chemical inertness, they
named argon, meaning lazy.
Besides argon, minute quantities of five other inert gases are
now known to be present in the atmosphere — helium (sun) , neon
(new) , krypton (hidden) , xenon (strange) , and radon (from ra-
dium) . The discovery and isolation of these gases from the air are
other amazing examples of precise and painstaking research. When
we consider that each of these gases is colorless, odorless, insoluble
in water, and chemically inert, refusing to unite with even the most
active elements, we begin to realize why they eluded chemists so long.
The inert gases go to work. For many years the inert gases re-
mained chemical curiosities. Helium, first identified in the sun in
1868 and later found in considerable quantities in natural gas, was
the first to be put to practical uses. It has taken the place of hydtogen
in inflating the balloons and blimps of the Army and Navy and some
104
NEW WORLD OF CHEMISTRY
of the weather observation balloons of the Weather Bureau. When
mixed with oxygen, it forms a synthetic air that is used under pres-
sure in caissons and is supplied to deep-sea divers to prevent the
bends.
The bends are a type of severe cramps caused by the sudden ex-
pansion and liberation of large quantities of nitrogen gas that have
entered the blood under the great pressures to which deep-sea divers
are subjected. When a diver rises to the surface where the air pres-
sure is considerably lower, less helium is capable of remaining dis-
solved in his blood, and, therefore, some of this gas is liberated from
the blood. Since helium is 40 percent less soluble in blood than is
nitrogen, less helium gas will be forced out of solution by the decrease
in pressure, and so the substitution of helium for the nitrogen of
common air helps to prevent the bends.
Laziness preferred. With the demand for lightweight metal parts
for airplanes, brought on by the tremendous need for aircraft in
World War II, research on the problem of welding magnesium was
stimulated and the helium-atmosphere process was perfected. In an
atmosphere of helium, oxidation of the magnesium cannot take
place, and the magnesium or magnesium alloys may be welded with
ease. The development of this method of welding magnesium is an
Linde Air Products Company
This workman is welding structural
aluminum by the sigma method. The
term tigma is derived from the first
letters of the words s/tiefaW inert gas
metal arc.
ATMOSPHERE 105
example of the way in which the needs of society and the research
of scientists are related.
Argon has replaced helium-arc welding for magnesium, and other
metals. It is also used to fill electric-light bulbs. When an evacuated
bulb is in use, metal evaporates from the filament, forming a deposit
on the inside of the bulb. This deposit blackens the bulb, making
it very inefficient. In an argon-filled bulb, this evaporation is re-
tarded and the lamp may be operated at higher temperatures than
if it had been evacuated only.
The ruby glow. The inert gases are also widely used in the glowing
glass tubes so familiar in advertising signs. When an electric current
is sent through a tube from which air has been removed and a minute
amount of neon gas introduced, the gas glows with an orange-red
light. The gas is at low pressure, about 12 mm. of mercury. The
amount of current required is extremely low, about ^ of an ampere,
but the voltage varies between 6000 and 12,000 volts. Neon and
krypton lights are also used to mark airplane routes and to signal
to airplane pilots. Small neon glow lamps are used in testing high-
frequency electric circuits such as those in radios.
When argon is used instead of neon, the light produced is blue.
However, most of the blue tubes of this type are filled with mercury
vapor rather than argon. Xenon gives a light blue light, and helium
a cream-colored or pale orange light. Following page 382 is an il-
lustration showing the colors produced by the inert gases in lighting.
The "idle" gases have thus been set to work. A leader in this field
predicted that, in time, much of our lighting would be done by
glowing gas in luminous tubes instead of by incandescent filaments.
Certainly it is true that much outdoor- lighting is now produced in
this way, but the prediction would not apply to indoor lighting in
which the field seems definitely being taken over by a newer develop-
ment, fluorescent tubes (see page 448) .
Radon gas, enclosed in sealed tubes, is used in the treatment of
cancer. The exact function of the inert gases in the air is still not
understood.
Our atmosphere also contains water vapor. Water vapor is always
present in air in varying amounts. The waters of the earth are con-
stantly evaporating. Plants give off immense quantities of water
vapor during transpiration, and animals, too, exhale water vapor.
Rain, dew, snow, fog, and other similar phenomena are caused by
the condensation of the invisible water vapor of the air. Frequently,
a pitcher of ice water sweats on the outside. This sweat is the water
that was formed when the water vapor of the air came in contact
106 NEW WORLD OF CHEMISTRY
with the cold outer surface of the pitcher and condensed. To de-
termine accurately the amount of water vapor in the air, we can
pass a known volume of air through a drying agent, such as calcium
chloride, CaCl2, or phosphorus pentoxide, P2O5. The increase in
weight in the drying agent equals the weight of the water vapor in
the sample.
CaCl2 + 2H2O -> CaCl2 2H2O
Carbon dioxide, too, is present in the air. Millions of tons of
carbon dioxide are poured into the air daily by the burning of
organic substances, by the decay of dead plant and animal matter,
and by the breathing of living things.
C + O2-»CO2
Carbon dioxide is a colorless, odorless gas fairly soluble in water,
and 1| times as heavy as air. Because carbon dioxide is already
completely oxidized, it does not burn. When passed through lime-
water (a water solution of calcium hydroxide) , it forms a white
precipitate, calcium carbonate, CaCO3. This formation of a white pre-
cipitate is the common test for carbon dioxide.
Ca(OH)2 + CO2 -» CaCO3 j + H2O
To determine accurately the amount of carbon dioxide present in
air, a known volume of air is passed through a concentrated solution
of potassium hydroxide, KOH, and the amount of potassium car-
bonate, K2CO3, that forms is determined.
2KOH + CO2 -» K2CO3 + H2O
Carbon dioxide and its uses are discussed in Chapter 23.
Cause and effect of atmospheric pollution. An adult inhales about
37 pounds of air a day, which is five times the weight of the food and
water that he consumes. We are very careful about getting pure
water, and have laws to protect us against the sale of impure food.
But we have not done as much about polluted air. In some cities
legislation has been passed to cut down pollution of the air caused
by smoke. Excessive smoke may impair health, damage crops, slowly
destroy property, and reduce visibility. Smog, a combination of
smoke and fog, is another serious problem (see pages 329, 616, and
617).
What is air conditioning? Not so many years ago, it was supposed
that the air in a crowded room was unhealthful because it contained
ATMOSPHERE 107
a large percentage of carbon dioxide and a lowered percentage of
oxygen. It has since been proved that even in a very crowded room
the percentage of carbon dioxide never reaches a point where it be-
comes harmful. The amount of oxygen rarely gets below 20 percent,
and it can be cut down even to 17 percent, at which point a candle
is extinguished, without being injurious to health.
Research has shown that "bad air" is really caused by high temper-
ature, lack of circulation, high percentage of water vapor, and various
odors that have accumulated.
Warm air can hold more water vapor than an equal volume of cold
air. A cubic meter of air at 20°C. (68°F.) , for example, is capable
of holding about 17 grams of water vapor, while the same volume of
air at 11°C. (52°F.) can hold only about one-half as much. The
ratio of the weight of water vapor present in air to the weight of
water vapor it is capable of holding under the existing conditions of
temperature and pressure is known' as its relative humidity.
The temperature of the human body is controlled in part by the
evaporation of perspiration. Evaporation absorbs heat. The amount
of heat absorbed depends upon the amount of perspiration evap-
orated. The cooling sensation produced depends upon both the
amount and the rate of evaporation.
Air with a high relative humidity, regardless of temperature, will
evaporate little perspiration. Consequently, on hot, damp days we
feel hotter than we do on hot days ihat are somewhat drier.
Tests have shown that in winter the air in many homes has a very
low relative humidity. Such very dry air has a twofold bad effect.
It tends to "dry us out." Evaporation of perspiration present proceeds
rapidly with great cooling effect. Consequently, the temperature of
dry air must be greater than the temperature of damper air for our
sensations of warmth to be the same. It t^kes more fuel to maintain
the higher temperature and, as a result, costs more. In addition,
temperatures above 21°C. (70°F.) are likely to produce drowsiness
and prevent us from doing our best work. Furthermore, such dry air
tends to dry out the linings of the nose, mouth, and throat, thus
causing great discomfort and reducing our resistance to common
colds and other respiratory diseases.
Today, steady progress is being made in supplying properly con-
ditioned air, not only to large auditoriums, factories, classrooms,
banks, office buildings, and railroad trains, but even to subway cars
and private homes. Air-conditioning equipment filters out dust and
pollen from the air, exhausts stale air, keeps the relative humidity
at the right point (about 50 percent) , and maintains a comfortable
108
NEW WORLD OF CHEMISTRY
TABLE 3. I
PERCENTAGE |
COMPOSITION I
OF DRY AIR I
temperature (about 68 °F.) . It also destroys the dulling quality of
"dead" air by keeping the air in motion. In addition to increasing
mental and physical efficiency, air conditioning is essential in certain
manufacturing processes, such as printing and the making of pills,
chocolate, rayon, paper, tobacco products, and steel bearings.
Summary: The composition of the atmosphere. The chief con-
stituents of samples of dry air taken near the surface of the earth are
shown in the following table.
1
CHIEF CONSTITUENTS OF DRY AIR
PERCEr
BY VOLUME
^TAGE
BY WEIGHT
Nitrogen
, Oxygen
Argon
Carbon dioxide
78,00
21.00
0.93
0.04
75.45
23.20
1.25
0.05
Air contains variable quantities of water vapor and carbon dioxide;
small amounts of the inert gases, argon, helium, krypton, neon,
radon, and xenon; and also minute amounts of other gases, such as
methane, carbon monoxide, hydrogen, nitrogen dioxide, and ozone,
as well as very finely divided solids, such as dust, bacteria, spores, and
pollen.
Liquid air. In Gullivers Travels, a famous Academy was visited,
and Jonathan Swift reports how some of its scientists were con-
densing air and letting the liquid flow like water. Probably Swift
believed the liquefaction of air a dream never to be realized. Yet
today liquid air is a common article of commerce and thousands of
tons of it are used every year.
The principles underling the manufacture of liquid air are:
(1) when a liquid evaporates, it absorbs heat from its surroundings
and thereby lowers their temperature; and (2) the sudden expansion
of a gas produces this same effect, When air is compressed, cooled,
and suddenly allowed to expand through a narrow opehing, its
temperature is lowered. If this process is repeated, a temperature is
finally reached which is low enough to liquefy the gas. A more de-
tailed treatment of the liquefaction of gases will be found on
page 253.
How the kinetic theory of matter explains the liquefaction of
air. All matter is thought to be made up of small particles (atoms
and molecules) that are in constant motion. In the case of gases, this
motion is extremely rapid (air molecules under normal conditions
ATMOSPHERE
109
move at about 20 miles per minute) . When any gas is cooled, its
molecules move more slowly, until finally a temperature is reached
at which the motion of the molecules of gas is so slow that they come
close enough together to form larger groups or clusters of molecules,
thus forming a liquid. When a gas expands, some heat is used in sep-
arating the molecules.
Compressing a gas has the effect of bringing the molecules closer
together. By a simultaneous cooling and compressing, any gas may
be liquefied. Some gases, such as chlorine, sulfur dioxide, and am-
monia, are easily liquefied. Other gases, such as oxygen, nitrogen,
hydrogen, and helium, require extremely low temperatures and
high pressures to change them to liquids. It has been calculated that
at — 273°C. (absolute zero) all motion of the particles of matter
ceases. The nearest approach to this temperature thus far attained
was made in 1952 when scientists at our National Bureau of Stand-
ards reached a temperature within 0.001,5° of absolute zero.*
* The measurement of absolute temperatures is discussed on p. 644.
Arthur D. Little, Inc.
Filling a flask with liquid helium
in an experimental laboratory.
Helium is liquefied by subjec-
ting it to both cooling and pres-
sure.
110
thermos
boffin
Dewar
flask
19. A thermos bottle
a Dewar flask. Note
imilarities. What is the
of the silvered sur-
faces?
Properties and uses of liquid air. Liquid air is a pale blue liquid
almost as heavy as water. It contains about 21 percent oxygen and
boils at — 190°C. When it evaporates, nitrogen boils off first and the
mixture becomes richer and richer in oxygen, for the same reason
that the boiling alcohol-water mixture in your car's radiator loses
alcohol faster than water, because the boiling point of alcohol is
lower than that of water.
Liquid air is used chiefly as a source of oxygen, nitrogen, and the
inert gases, which boil off at different temperatures. The Nazis used
liquid oxygen and alcohol to fuel the V-2 rocket bombs which they
hurled against England in 1945.
Because of the great tendency of liquid air to evaporate, small
volumes of it are kept in special containers called Dewar flasks, which
are similar in construction to the familiar thermos bottle. These
flasks cannot be tightly stoppered, for any attempt to confine liquid
air too closely results in explosion. Because of the danger of injury
both from such explosions and from contact with a substance at such
an extremely low temperature, persons handling liquid air must use
great care.
Air Reduction Company, Inc.
Inferior of an oxygen-nitrogen
plant. The workman is standing in
front of a column in which the two
gases are obtained from liquid air.
ATMOSPHERE 111
The properties of substances change when immersed in liquid air.
Liquid mercury becomes solid enough to be used as a hammer head.
Rubber turns hard and brittle. The resistance of copper to an elec-
tric current is decreased 50 times.
YOU WILL ENJOY READING
Cady, H. P. "Liquid Air." Journal of Chemical Education,
June, 1931, pp. 1027-1043. A fascinating account of experi-
ments with liquid air.
Kaempffert, Waldemar. Explorations in Science, Chapter 7,
pp. 102-108 is entitled, "This most excellent canopy, the air."
The Viking Press, New York, 1953.
Ramsay, William. The Gases of the Atmosphere, pp. 148-
181, 234-269. The Macmillan Co., London, 1915. Sir William
tells of the discovery of argon and discusses other inert gases.
USEFUL IDEAS DEVELOPED
1. Air is a mixture of gases because (1) it has no definite
composition, (2) when air dissolves in water, the dissolved
air contains more oxygen than common air, and (3) liquid air
does not have a definite boiling point.
2. Air conditioners cleanse the air, keep the relative humid-
ity where it belongs (about 50 percent) , maintain a comfort-
able temperature (68 °F.) , and keep the air in motion.
3. The kinetic theory of matter assumes that gases are made
up of small particles (atoms and molecules) in active motion.
Cooling the gas slows down this motion and brings the par-
ticles closer together until a liquid is formed. The molecules
of liquids and solids are also in motion.
USING WHAT YOU HAVE LEARNED
Group A
1. Who first definitely proved that air has weight?
2. (a) What pressure does the atmosphere exert at sea
level? (b) What causes this pressure?
3. What are the differences between the troposphere, strat-
osphere, and Heaviside layer?
4. Does the composition of the atmosphere prove that air
is a mixture? Explain.
5. Give two reasons other than composition for believing
that air is not a compound.
112 NEW WORLD OF CHEMISTRY
6. State the composition of dry air by volume.
7. What substances are found in air in variable quantities?
8. Make a diagram illustrating a laboratory method for
determining the percentage of O2 in air.
9. (a) What is the gas that is left in the measuring tube
used in determining the percentage of O0 in air? (b) What
are the impurities in this gas?
10. State four physical properties of N2.
11. What is the chief chemical property of N2?
12. Why is it wrong to call N2 an inert gas?
13. Write a chemical equation illustrating the action of
hot Mg on N2.
14. Under what conditions does N., combine chemically
with O3 and H2? »
1 I
15. (a) State two functions of the N2 of air. (b) State two
commercial uses of N2.
16. (a) Why is Mg difficult to weld? (b) Under what con-
dition is it easily welded? (c) Development of this process
is an example of what?
17. Name the six inert gases of the atmosphere.
18. What five properties are common to all six of the inert
gases of the atmosphere?
19. Match each element listed in the first column with the
correct item in the second column.
1) Ne a) hidden
2) A b) balloons
3) Kr c) strange
4) He d) luminous tubes
5) Xe e) electric-light bulbs
f) airplanes
20. Explain the "sweating" of a pitcher filled with ice water,
21. How would you determine the amount of water vapor
present in air?
22. How would you determine the amount of CO, present
in air?
23. State (a) four physical and (b) two chemical properties
of C02.
24. What is air conditioning?
25. Name (a) two causes and (b) three effects of atmos-
pheric pollution.
ATMOSPHERE 113
26. What two conditions in the air of a poorly ventilated
room make it both uncomfortable and unhealthful?
t . . .
27. What is relative humidity?
28. Upon what principle does the liquefaction of a gas
depend?
29. Discuss the meaning of the kinetic theory of matter.
30. What is the chiet use of liquid air?
31. State three properties of liquid air.
32. Compare the effectiveness of determining the percentage
of O2 in air by the phosphorus and the iron filings methods.
Group B
33. Why does the composition of air by weight differ from
its composition by volume?
34. A piece of burning charcoal is plunged into liquid air.
It keeps on burning with even greater splendor. Why does not
the extreme cold of liquid air extinguish the burning charcoal?
35. Explain the presence of vast amounts of free N2 in air,
and only relatively small amounts of nitrogen compounds in
the earth's crust.
36. In different samples of air the following substances are
placed: CaCl,, P, hot Mg, and Ca (OH) ,. Explain what hap-
pens in each case.
37. Why does a kettle of liquid air boil when placed on ice?
38. "Coal burned in our furnace returns to us in our
bread." Explain.
39. What weight of dry air would be theoretically needed
to extract ten grams of pure oxygen?
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Visit a neon-sign factory or an air-conditioned building
or factory and make a full report on your visit to the class.
2. Prepare simple demonstrations to show that (1) air
contains water vapor, and (2) your breath contains CO2. Per-
form these experiments before your class the following day.
3. With the aid of a wet and dry bulb thermometer, de-
termine the relative humidity of your classroom.
4. Prepare a report on the probable origin of the earth's
atmosphere. Use a good book on physiography or one of the
references listed.
8
EQUATIONS:
SHORTHAND
OF CHEMISTRY
. . . Here evidently we are at the
birthplace of the chemical equation,
yet we cannot find in the writings of
Lavoisier this instrument as we know
it; for our chemical equation de-
pends as much on the atomic theory
as on the doctrine of the conservation
of mass. Caven and Cranston, 1928
Chemical equations, the shorthand of chemistry. Chemists use
symbols and formulas that are understood to represent elements and
compounds by scientists in all parts of the world. But chemists make
further use of these symbols and formulas. They use them to tell the
story of chemical change — the reacting substances, the type of chem-
ical change, the products of the reaction, and various other facts.
All of these facts are expressed in the form of a chemical equation,
which, in certain respects, is similar to the equations you have used
in arithmetic and algebra.
Considering the chemical reaction that takes place when iron and
sulfur are heated together, chemists and students of chemistry write
as follows:
Fc + S -» FcS
As you see, the two sides of the equation are separated by an arrow,
-». This arrow means yields, or forms, and should be read using
these or similar words. On the left side of the arrow, chemists write
the symbols or formulas of the substances that react. On the right side
of the arrow, they write the symbols and formulas of the products of
the reaction. A word-equation that will express the same reaction
as the chemical equation that we have just been discussing follows:
114
EQUATIONS
115
Iron reacts with sulfur to yield iron sulfide.
Equations represent reality. To write a chemical equation, we
must know the formulas of the substances involved. To write a
chemical equation correctly, we must know exactly what substances
are reacting and what substances are the products of the reaction.
The facts are determined before an equation is written.
We can write equations for only those chemical reactions known to
be capable of actually taking place. Chemical equations must repre-
sent actual conditions. For example, it is incorrect to write He + S
-» HeS, because helium and sulfur do not react together, and
helium sulfide, HeS, has never been prepared.
Balancing an equation. Lavoisier once wrote: "One may take it
-for granted that in every reaction there is an equal quantity of matter
before and after the operation." Atoms do not disappear in the
process of chemical reaction. Therefore, the number of atoms in
the reacting substances must equal the number of atoms in the
products of the reaction. The number of atoms of each element must
be the same on each side of the arrow. Every correct equation con-
forms, in this way, to the law of the conservation of matter.
Points to remember in balancing equations.
1) Elements can occur in a free state, not combined with other
elements. For reasons which you will learn later (pages 279—280) ,
a number of these elements, including oxygen, hydrogen, nitrogen,
chlorine, bromine, and iodine occur as molecules containing 2 atoms.
They are written O,, H.,. N2, C12, Br2, and I2, thus:
H2 + C12 — » 2HC1 (hydrogen chloride)
N2 4- O2 — > 2NO (nitric oxide)
. . < Cu 4- Br2 — > CuBr2 (copper bromide)
2H2O -> 2H2 + 02
Standard Oil Company
Delicate measurements must be
made to determine the actual
conditions that are described
by chemical equations.
116 NEW WORLD OF CHEMISTRY
2) In general, a radical remains unaltered during a chemical
change and its symbols are carried over to the right side of the equa-
tion unchanged:
Zn + H2SO4 -> ZnSO4 + H2
2Na + 2HOH -» 2NaOH + H2
BaO2 + H2SO4 -> BaSO4 + H2O2
Balancing an equation consists of changing coefficients until the same
total number of atoms of each element is shown on each side of the
arrow.
Do not alter the subscript of a radical nor any other subscript in
order to make an equation balance. Such alteration would mean a
change in the actual composition of the compound. This in turn
would mean an entirely different substance which does not actually
appear in the chemical change. We may, however, alter coefficients
of any element or compound without changing the meaning of the
chemical symbols. The coefficient describes the number of molecules
of a substance. By changing the number of molecules present in a
chemical change, we do not alter the composition of the substances
which are involved.
Your best preparation for writing equations correctly is a thorough
knowledge of valence and of formulas. In addition, you should
attack the problem in a systematic and thoughtful way. The follow-
ing procedure should be helpful.
HOW TO BALANCE EQUATIONS
1) Write the equation without giving coefficients to any of
the formulas.
2) Write any free element occurring in the equation without
a subscript. Retain all other subscripts.
3) Select the compound with the greatest number of atoms.
For one of its elements, compare the number of atoms appear-
ing on each side of the equation. If the numbers differ, decide
upon a coefficient or coefficients which will equalize the num-
ber of atoms on each side.
4) Repeat this procedure for the other elements until the
equation is balanced, that is, until the same number of atoms
of each element appears on each side of the equation.
5) Add the subscript to any free element in the equation
which may require it. At this stage, free elements may have
been given coefficients. If the coefficient of the free element is
EQUATIONS 117
an even number divide the coefficient by two when the sub-
script is added. The total number of atoms of the free element
will then remain the same (for example, 6H contains the same
number of atoms as 3H2) .
6) If, however, the coefficient of the free element is an odd
number, it cannot be divided by two. In this case, when the
subscript is added to the free element, all other formulas on
both sides of the equation must be multiplied by two.
EXAMPLE A: Write the equation for the reaction which takes
place when potassium chlorate is heated to form potassium
chloride and oxygen.
1) Write the equation without coefficients. Note that the
free oxygen which is usually written O2 is written at this stage
without a subscript.
KClOa -^ KC1 + O
2) KC1O3 is the compound with the largest number of
atoms. By inspection, we see that oxygen is the only element
in the compound which does not appear with an equal number
of atoms on the other side of the arrow. To make the O on the
right balance the O3 on the left, we multiply O by the coeffi-
cient 3:
KC1O3 -> KC1 + 3O
3) As we have just learned, free oxygen must be written as
•O.,. Add the subscript to the free oxygen on the right of the
arrow. The coefficient 3 is not divisible by two, so when we
add the subscript, we find there are six atoms of free oxygen
on the right side of the arrow, but only three atoms of oxygen
on the left side:
KC1O3 -> KC1 + 3O2
4) We may make the equation balance again by multiplying
.all other formulas by two:
2KC1O3 -> 2KC1 + 3O2
EXAMPLE B: Write the equation for the burning of benzene
(C6H0) to form CO2 and H2O.
1) Write the unbalanced equation. Do not write the sub-
.script for the free oxygen.
O -> CO2 + H2O
118 NEW WORLD OF CHEMISTRY
2) C6H6 is the compound with the largest number of atoms.
It contains 6 atoms of carbon, but there is only one atom of
carbon to the right of the arrow. Therefore, we place the
coefficient 6 in front of CO.
3) C6H6 also contains 6 atoms of hydrogen, but there are
only 2 atoms of hydrogen on the right side of the arrow. There-
fore, we place the coefficient 3 in front of H2O.
CeH6 + O -> 6CO2 + 3H2O
4) We now see that there are 15 atoms (6 X 2 + 3 X *) of
oxygen on the right, but only one atom of oxygen on the left.
Therefore, we place the coefficient 15 in front of O.
C6Hfl + 150 -> 6C02 + 3H2O
5) Since free oxygen is written O2, we add the subscript.
6) Since 15 is not divisible by two, we may bring the equa-
tion into balance once again by multiplying all other formulas
by the coefficient 2.
SQHe + 15O2 -> 12CO2 + 6H2O
If the coefficient of the oxygen had been an even number, we
could have divided it by two when adding the subscript. It
would then be unnecessary to multiply the other formulas by
the coefficient 2.
HOW TO CHECK EQUATIONS. To check an equation, proceed as follows:
1) Examine the first symbol on the left. Compare the num-
ber of atoms against the number of atoms of the same element
on the other side of the equation. If the numbers are equal,
put a small check over the symbols for the element.
2) Proceed to each symbol in turn until you have put a check
over every symbol in the equation.
Thus, in checking the final equation in Example A above,
2KC1O8 -* 2KC1 + 3O2, examine the first symbol, K. There are two
atoms of potassium on each side. Put a check over the K on each side.
Examine the next symbol, Cl. Again we have two atoms on each
side. Check them. Finally, we have six atoms of oxygen on each side,
which we check. All the symbols are then checked and accounted for:
v/ v/ v/ V \/ v/
2KC1O3 -> 2KC1 + 3O2
EQUATIONS
119
The equation is now balanced accurately. It is correct both mathe-
matically and chemically. Following the same procedure, check the
final equation in Example B.
Fig. 20. The of th« tuft contains o
of oncf an test holding a
of AgN03. When the |t inverted
(right), A§CI is but the cca|e tho
no change. No matter hot
destroyed.
2.
AgCI
PRACTICE WORK ON BALANCING EQUATIONS
To be sure that you understand the suggestions just given for bal-
ancing equations, balance the following:
a) NaClOs -> NaCl + O g)
b) Hg + 0-»HgO h)
c) P205 + H20 -> HPO, i)
d) Mg + 0->MgO j)
e) Na + Cl->NaCl k)
f) C + O->CO
Gu + O -> CuO
Fe -f O — > Fe2O3
O + P -» P20»
P 4- O -> P2O5
SiO2 -f G -* SiC
CO
What are the four types of chemical reactions? There are four
general types of chemical reactions.
1) Direct combination, or synthesis. When two or more elements
or compounds combine directly, forming a chemically more complex
122 NEW WORLD OF CHEMISTRY
t . . .
6. Why must subscripts never be changed in balancing
equations?
7. What is wrong with the following:
2Ne (neon) + O2 -» 2NeO?
8. Why may coefficients be changed in balancing an equa-
tion?
9. Why is a knowledge of valence so essential in writing
equations?
10. Name five elements which, when free, must be written
with the subscript 2. ft
a 1
11. Correct this: Mg -(- Br -» MgBr2.
12. Generally, what happens to a radical during a chemical
change?
13. Balance the following equation, giving your reason for
each step:
Acetylene (C2H2) -f oxygen (O2)
-> carbon dioxide (CO2) -\- water (H2O)
14. What are the four general types of reactions?
15. Give an example of an equation illustrating direct
combination.
16. Give an example of an equation representing simple
decomposition.
17. How do you recognize an equation representing simple
replacement?
18. Give an example of an equation representing double
replacement.
19. What is the general rule followed in balancing a double-
replacement equation?
20. Complete and balance the following equation. Show
how to check it to see if it is correctly balanced.
Na -f HOH -» NaOH + H
21. Show in detail how to balance the equation represent-
ing the decomposition of sodium chlorate, NaClO3.
22. When CaCl2-2H2O is heated, water is liberated. Write
the equation.
EQUATIONS
23. Balance the following equations. Check each one. Do
not write in this book.
a) Cu + S->Cu2S
b) P + C12->PC13
c) C + C02->CO
d) H2S + Pb(NO3)2 -> HN03 + PbS
e) CaO + G -> GaC2 + CO
f) SiO2 + G -» SiC + CO
24. Complete and balance the following equations. Check
each one. Do not write in this book.
a) H2O2 — > e) BaO2 + H2SO4 — >
b) HgO-> f) AgNO3 + KCl-»
c) Fe + HCl-> g) Mg + H2S04-»
d) CuO + H2-» h) Ca(OH)2 + CO2 ->
25. Write balanced equations for the following. Name each
element and compound that appears in the equation. Check
each equation.
a) The electrolysis of H2O.
b) The laboratory preparation of O2.
c) The equation representing Priestley's discovery of OL>.
d) The laboratory preparation of H2.
e) The heating of Cu in air.
f) The preparation of H2O2.
g) The reduction of CuO^ by H2.
h) The heating of crystallized washing soda,
i) The chemical reaction of the oxy hydrogen torch,
j) The action of Na on H2O.
k) The passing of N2 over hot Mg.
1) The reaction between H2O and CO2.
26. Is there anything in a balanced equation which did not
have to be determined by experimentation?
MATHEMATICS
OF CHEMISTRY
. . . As the usefulness and accuracy
of chemistry depend entirely upon
the determination of the weights of
the ingredients and products, too
much precision cannot be employed
in this part of the subject, and for
this purpose we must be provided
with good instruments.
Lavoisier, 1743-1794
Measurement in chemistry. Chemists are chiefly concerned in
analyzing substances or in making new substances. One question
they always ask is how much of each element is present in a sub-
stance. Their measurements of quantity must be very exact. Some-
times their answer is in terms of the volume occupied by an element.
More often the answer they want is in terms of weight.
Here in America, in everyday affairs, we measure the weight of
substances in ounces and pounds. Chemists use a different kind of
measurement. They weigh substances in grams, kilograms, and micro-
grams. They also measure the weight of atoms in terms of atomic
weight.
All measures of weight are comparisons. When you say, "My
friend, Charlie Ross, weighs 132 pounds," you are making a definite
statement that anyone can understand. But what is a pound? You
know what a pound of butter looks like, but a pound by itself
doesn't look like anything.
Down in the Bureau of Standards in Washington is a carefully
protected metal cylinder called the prototype, or standard, kilogram.
Every measurement of weight made in our nation is a comparison
with that particular kilogram. For example, the pound which we
124
MATHEMATICS OF CHEMISTRY 125
use in everyday measurements of weight may be defined as 0.4536
of the standard kilogram.
Just as the pound is divided into ounces, the kilogram may be
divided into smaller units known as grams. One thousand grams
equal one kilogram. The gram is a customary unit of weight in chem-
istry. When chemists speak of ten grams of iron, they are actually
referring to an amount weighing T^ of the standard kilogram in
Washington.
Chemists also use a Table of Atomic Weights in which one atom of
each element has a particular weight. The weight assigned to each
element, like all measures of weight, is a comparison. How was this
Table of Atomic Weights made?
Dalton determines atomic weights. As you know, Dal ton believed
that the atoms of different elements have different weights. Dalton
knew that if he could find the weights of the atoms, the progress of
chemistry would be speeded. He rcali/.ed that he could not actually
weigh an atom of an element. In fact, it took more than 100 years
from the time of Dalton's experiments until accurate methods and
precise instruments made it possible to determine the actual weights
of atoms.
However, Dalton knew that elements combine according to fixed
ratios by weight. For example, 22.997 grams of sodium combine with
79.916 grams of bromine to form sodium bromide — a ratio of ap-
proximately 1 : 3.5. The two elements never vary in the ratio of their
weights in forming this compound. In a similar manner, all elements
combine according to certain specific ratios by weight. Dalton be-
lieved that the ratios of these combining weights depended upon the
weights of the atoms of each element. He believed further that, by
studying the ratios of the weights in which elements combine, he
would be able to determine the relative weights of the atoms of the
elements.
The prototype kilogram No. 20
is a platinum-iridium cylinder
39 mm. in diameter and 39 mm.
high. It is kept at Washington,
D.C., in the laboratories of the
Bureau of Standards.
126 NEW WORLD OF CHEMISTRY
He selected the atom of hydrogen for his standard and assigned to
it the atomic weight 1. By choosing the lightest element for a stand-
ard, he made sure that the atomic weight of each of the other ele-
ments would be greater than one. Then in a series of experiments,
he found how the weight of the atoms of 13 other elements com-
pared with his standard. He found that an atom of oxygen weighed
seven times as much as an atom of hydrogen and assigned to oxygen
the atomic weight of 7; he found that an atom of phosphorus
weighed nine times as much as an atom of hydrogen and assigned to
phosphorus the atomic weight of 9. These weights have since been
found in error, but that in no way detracts from the value of Dalton's
accomplishment.
Dalton prepares an historic table. From his experimental data,
Dal ton prepared a list of the 14 elements arranged according to the
increasing relative weight of their atoms. This was the first table of
atomic weights. It was first made public on October 21, 1803 "before
a select group of nine members and friends in the rooms of the
Literary and Philosophical Society of Manchester." Although inac-
curate, the table compiled by this Quaker schoolteacher remains a
monument to his foresight and intellectual accomplishment. His
achievement was a crucial advance in chemistry and formed the
keystone of his theory.
What is the atomic weight of an element? Later investigators fol-
lowed Dalton's method, but used oxygen as a standard rather than
hydrogen. They found oxygen a better choice because it combines
with far more elements than hydrogen. The atomic weight of oxygen
was given the whole number 16. Many of the other elements have
atomic weights that are whole numbers. The weight of hydrogen
remains approximately one — actually 1.0078 in comparison to
oxygen.
The modern Table of Atomic Weights appears on the opposite
page. For practical purposes, since the atomic weight of hydrogen is
approximately one, hydrogen can still be used as the basis of com-
parison. The number 200.6 after mercury means that one atom of
mercury weighs a little less than 200 times as much as an atom of
hydrogen. It weighs a little more than 12| times as much as an atom
of oxygen.
The atomic weight of an element is a number that shows the
comparison of the weight of one of its atoms to the weight of one
atom of oxygen, which is considered to be 16.
Always remember that atomic weights are not measured weights
like an ounce or a gram. They are merely relative weights. The
TABLE 4. APPROXIMATE ATOMIC WEIGHTS
ELEMENT
SYMBOL
APPROX.
AT. WT.
ELEMENT
SYMBOL
* APPROX.
AT. WT.
Aluminum
Al
27
Lead
Pb
207
Antimony
Sb
121.8
Lithium
Li
7
Arsenic
As
75
Magnesium
Mg
24.3
Barium
Ba
137.4
Manganese
Mn
55
Bismuth
Bi
209
Mercury
Hg
200.6
Boron
B
11
Nickel
Ni
58.7
Bromine
Br
80
Nitrogen
N
14
Cadmium
Cd
112.4
Oxygen
O
16
Calcium
Ca
40
Phosphorus
P
31
Carbon
C
12
Platinum
Pt
195
Chlorine
Cl
35.5
Potassium
K
39
Chromium
Cr
52
Radium
Ra
226
Cobalt
Co
59
Silicon
Si
28
Copper
Cu
63.5
Silver
Ag
108
Fluorine
F
19
Sodium
Na
23
Gold
Au
197
Strontium
Sr
87.6
Hydrogen
H
1
Sulfur
S
32
Iodine
1
127
Tin
Sn
118.7
Iron
Fe
56
Zinc
Zn
65.4
actual measured weight of each of the atoms is an extremely small
quantity, inconvenient to use in most calculations. For example,
the actual weight of an atom of oxygen is about 0.000,000,000,000,-
000,000,000,026 gram.
The Table of Atomic Weights is the foundation of chemical
mathematics. Realizing the importance of accurate atomic weights,
many chemists performed a tremendous number of experiments to
make the Table of Atomic Weights as free from error as possible.
Chemists of all countries cooperated in this huge undertaking. In
our own country, Theodore W. Richards and his students at Harvard
spent almost half a century in this epoch-making work. National
boundaries were forgotten and men from all over the world pooled
the results of their experiments to give us our present Table of
Atomic Weights. Research chemists, industrial chemists, and student
chemists all depend upon the Table of Atomic Weights in making
their chemical calculations.
Solving type problems in chemistry. Most of the problems met in
elementary chemistry can be grouped conveniently into five types.
The first two, and one variety of the third are described below. Two
additipnal varieties of the third type are discussed in Chapter 19
and two more types, more complex and less frequently met, are
discussed on pages 636—641.
127
128 NEW WORLD OF CHEMISTRY
With a thorough understanding of the type problems discussed
in this book, you should have no trouble in solving practically all
the common chemical problems. Frequent reflective practice is, of
course, necessary for mastery. Hence problems of various types are
included in the questions at the end of each of the remaining
chapters.
The meaning of symbols. A knowledge of the meaning of chemical
symbols is essential in solving the various types of chemical problems.
The symbol for an element, like K, does three different jobs. First,
it may be used to name the element. Second, it may be used to mean
one atom of that element. Third, it may stand for one atomic weight
of the element. Thus K is the symbol for potassium, for one atom of
potassium, and for 39 units of potassium in any system of measuring
weight (as 39 grams, 39 ounces, 39 pounds) .
The chemical symbols for a formula like KCl also do three jobs.
First, they represent the name of the compound, potassium chloride.
Second, they represent one molecule of potassium chloride. Third,
they represent one molecular weight of the compound. All of the
type problems involve the use of both atomic iveight and molecular
weight.
TYPES OF PROBLEMS
TYPE 1: TO FIND THE MOLECULAR WEIGHT* OF A COMPOUND
FROM ITS FORMULA
Molecular weight is the ratio of the weight of one molecule
of a compound to the atomic weight of oxygen (16) . Like
atomic weight, it is only a relative weight. The molecular
weight of a compound is obtained by adding together the
atomic weights of each of the atoms in one molecule of the
compound.
Procedure.
1. Find the atomic weights of the elements in the chart on
page 127. Place these numbers under the symbols and add
them.
* Some compounds do not exist as molecules and therefore cannot have a
molecular weight. However, the formulas of these compounds are written in the
conventional manner. The steps given in this discussion for finding molecular
weight may also be used for finding the formula weight of the non-molecular
compounds. For the purposes of this book, formula weight may be considered the
same as molecular weight.
MATHEMATICS OF CHEMISTRY 129
EXAMPLE: Find the molecular weight of potassium chloride,.
KCL
K a
39 + 35.5 = 74.5, the mol. wt. of KC1
2. When an element is followed by a subscript, be sure to
multiply the atomic weight of the element by the subscript.
EXAMPLE: Find the molecular weight of magnesium sulfate,
MgS04.
Mg S 04
24 + 32 + (16 X 4) = 24 + 32 + 64
= 120, the mol. wt. of MgSO4
3. When a formula is preceded by a numerical coefficient,
multiply the total molecular weight by the coefficient to find
the relative weight (rel. wt.) .
EXAMPLE: Find the relative weight of two molecules of mag-
nesium carbonate, 2MgCO3.
2(Mg C 08)
2 [24 + 12 + (16 X 3)]
2 (24 + 12 + 48)
2 (84) = 168, the rel. wt. of 2 molecules of MgCO3
4. When a radical is enclosed in parentheses followed by a
subscript, multiply the sum of the atomic weights of all the
elements of the radical by the subscript.
EXAMPLE: Find the molecular weight of calcium bicarbonate,
Ca (HC03)2.
Ca (H C O8)2
40 + [1 + 12 + (16 X 3)]2
40 + (1 + 12 + 48)2
40 +(6 1)2
40 + 122 = 162, the mol. wt. of Ca (HCO8)2
5. Water of crystallization (see page 68) is chemically part
of certain compounds and is separated from the rest of the
formula by a dot. The dot stands for plus and is not to be
considered a multiplication sign.
130 NEW WORLD OF CHEMISTRY
EXAMPLE: Find the molecular weight of crystallized copper
sulfate, CuSO4 • 5H2O.
Cu S 04 • 5H2 O
64 + 32 + (16 X 4) + 5(1 X 2 + 16)
64 + 32 + 64 +5(18)
64 + 32 + 64 + 90 = 250, the mol. wt. of crystallized copper
sulfate.
The gram-molecular weight or mole. A very convenient unit in
many calculations is the gram-molecular weight or mole, which is
used in chemical equations (see pages 282-283) and in preparing
standard solutions (see pages 207-208) . A mole is the molecular
weight of a substance expressed in grams. For example, a mole of
potassium chloride (see Procedure 1) is 74.5 grams.
PRACTICE WORK ON PROBLEMS OF TYPE 1
1. Calculate the molecular weight of (a) KBr and (b) Nal.
2. Find the molecular weights of (a) LiCl and (b) ZnO.
3. What is the weight of a mole of K3PO4?
4. Find the molecular weights of (a) H2SO4, (b) CaCO3>
and (c) BaSO3.
5. Calculate the molecular weights of (a) Cu (HCO3) 0,
(b)Ba(N03)2, and (c) A12 (SO4) ,.
6. Calculate the molecular weight of Na2S2O3 • 5H2O (com-
monly called hypo) .
7. What is the weight of a mole of gypsum, CaSO4 • 2H2O?
8. Find the molecular weight of plaster of Paris,
(CaS04)2.H20.
9. Thfe formula of glauber salt is Na2SO4 • 10H.,O. Find
the relative weight of three molecules of this substance.
10. Calculate the relative weight of 4Na2B4O7 • 10H2O.
TYPE 2: TO FIND THE PERCENTAGE COMPOSITION OF A COM-
POUND FROM ITS FORMULA
The percentage composition of a compound is found by
computing the percentage by weight of each different element
in the compound. This is a simple percentage problem.
Procedure. Divide the atomic weight of each element by the molecu-
lar weight of the compound and multiply the fraction thus
obtained by 100.
MATHEMATICS OF CHEMISTRY 131
EXAMPLE A: Find the percentage composition of nitric acid,
HNO8.
H N O3
1-1- 14 + (16X3)
1 + 14 + 48 = 63, the mol. wt. of HNO3
% of hydrogen
at.wt.ofH 1 X 100
mol.wtofHNO, ~ 63
% of nitrogen
x 100 - 22.2%
mol. wt. of HNO3 --- 63
% of oxygen
_ rel.wt.of30 48X100
X
mol. wt. of HN03 63
Total = 100.0%
EXAMPLE B: Find the percentage of water of crystallization in
BaCl2 - 2H2O.
Ba C12 2H2 O
137 + (35.5 X 2) + 2(1 X 2 + 16)
137 + 71 +2(18)
137 + 71 + 36 = 244, the mol. wt. of BaCl2 2H2O.
Percentage = rcl. wt. of 2H2O x 1Q()
of water mol. wt. of crystal
EXAMPLE C: Find the percentage composition of
Na2SO4 10H2O.
Na2 S O4 • 10H2 O
(23 X 2) + 32 + (16 X 4) + 10(1 X 2 + 16)
46 + 32 + 64 + 180 = 322, the mol. wt. of the compound.
Now find the total atomic weight of each element in the
compound, thus:
Sodium - 2 atoms - 23 X 2 = 46, the total at. wt. of Na
Sulfur = 1 atom = 32, the total at. wt. of S
Oxygen = 14 atoms = 16 X 14 = 224, the total at. wt. of O
Hydrogen • 20 atoms = 20 X 1 = 20, the total at. wt. of H
Sum of the at. wt. = 322, the mol. wt. of the
compound.
132 NEW WORLD OF CHEMISTRY
% of sodium = -3 X 100 = 14.3%
% of sulfur = ^ X 100 = 10.0%
% of oxygen - ||| X 100 - 69.5%
% of hydrogen = ~ X 100 = 6.2%
Total = 100.0%
EXAMPLE D: Find the weight of iron in 80 Ib. of an ore con-
taining 90 percent ferric oxide, Fe2O3.
The weight of ferric oxide in the ore is 90 percent of 80 Ib.,
or 72 Ib.
Fe2 O3
(56 X 2) + (16 X 3) = 160 = mol. wt. of Fe2O3
Percentage of iron in Fe2O3 = H$ X 100 = 70%
Therefore 70% of 72 Ib., or 50.4 Ib., is the weight of the iron
in the ore.
PRACTICE WORK ON PROBLEMS OF TYPE 2
1. Calculate the percentage composition of (a) water,
(b) hydrogen peroxide, and (c) mercuric oxide.
2. Calculate the percentage composition of (a) H2CO3,
(b) N2O4, and (c) CaSO4 - 2H2O.
3. Find the percentage composition of chrome alum, *
KCr(S04)2- 12H20.
4. Calculate the percentage composition of crystallized potas-
sium ferrocyanide, whose formula is K4Fe (CN) 6 • 3H2O.
5. Find the percentage of oxygen in a compound whose
formula is NiSO4 - (NH4) 2SO4 - 6H2O.
6. How much aluminum can be obtained from 100 Ib. of its
cryolite ore which, upon analysis, shows the presence of 80
percent Na3AlF0?
TYPE 3: PROBLEMS BASED ON CHEMICAL EQUATIONS
Because the symbol of an element and the formula of a com-
pound may represent definite weights, an equation also may
be considered to represent definite weights of the substances
taking part in the reaction. Thus, 2Ag -f- S — » Ag2S may be
read, 216 grams of silver, plus 32 grams of sulfur yield 248
grams of silver sulfide. Note that the actual weights are based
on the atomic weights.
Problems based on chemical equations may be broadly di-
vided into three groups: A. Straight- weight problems;
B. Weight-volume problems; and C. Straight- volume problems.
MATHEMATICS OF CHEMISTRY 133
Group A is discussed below. Groups B and C are discussed
later in the book.
TYPE 3A: STRAIGHT-WEIGHT PROBLEMS
Straight-weight problems involve finding one weight in an
equation when another is given.
EXAMPLE: How many grams of calcium carbonate will be
formed by the complete reaction between 222 g. of calcium
hydroxide and carbon dioxide?
Procedure.
1. Write the balanced equation.
Ca(OH)2 + C02 -» H20 + CaCO3
2. Write the given weight over its formula. Write x over the
formula of the substance whose weight is to be found. Cross
out all other formulas in the equation.
222 g. x g.
Ca(OH)2 +^Q*,^H2QL+ CaCO8
3. Since the same relationship exists between the actual
weights as exists between the molecular weights represented
in the equation, write the molecular weights of the substances
involved under their respective formulas. Do not ignore any
coefficient.
222 g. x g.
Ca (O H)2 -» Ca C O3
40 + (16 + 1)2 40 + 12 + (16 X 3)
40 + (17X2) 40+ 12 + 48
74 100
4. Write the mathematical equation represented by the
known and unknown weights. Solve for x.
222 = _£_
74 100
74* = 22,200
x = 300, the number of grams of CaCOs produced
Alternate method. We can avoid the use of an equation in-
volving x by solving the problem as follows:
134
wt. of substance used
NEW WORLD OF CHEMISTRY
mol. wt. of sub-
X P - = Answer
mol. wt. of substance used stance formed
™ X 100 = 300, the number of grams of CaCO3 produced.
PRACTICE WORK ON PROBLEMS OF TYPE 3A
1. How much magnesium is required to react with suffi-
cient hydrochloric acid to produce 10 g. of hydrogen?
Mg + 2HC1 -* MgCl2 + H2
2. 434 g. of mercuric oxide, HgO, were decomposed by
heat. How much mercury was formed?
3. How much potassium chlorate would be needed to
prepare 384 g. of oxygen?
4. How much hydrogen would be needed to reduce com-
pletely 100 g. of cupric oxide, CuO?
5. 11.5 g. of sodium react completely with water. What
weight of sodium hydroxide is formed?
6. By the electrolysis of water 12 Ib. of hydrogen were
liberated. What weight of oxygen was formed at the same
time?
7. (a) What weight of magnesium will be needed to react
with sulfuric acid to produce 30 g. of MgSO4? (b) What
weight of hydrogen will be evolved?
8. After heating, 10 g. of crystalli/ed copper sulfate gave
6.4 g. of the anhydrous salt, CuSO4. Calculate the number of
molecules of water of crystallization in the original compound.
Let x represent the number of molecules of water of crystal-
lization.
10 g.
CuSO4-*H2O
6.4 g.
GuSO4
3.6 g. (that is, 10 - 6.4)
Now complete the problem.
Standard Oil Company (A'./.)
Delicate instruments such
as these are used in micro-
chemistry, the branch of
chemistry which involves
handling extremely small
quantities of matter.
MATHEMATICS OF CHEMISTRY 135
9. After being heated, 10 g. of crystallized washing soda
gave 3.71 g. of Na2CO3. Calculate the number of molecules
of water of crystallization.
10. If 4 g. of crystallized barium chloride lost 0.59 g. upon
being heated to constant weight, find the formula of the
crystalline salt.
YOU WILL ENJOY READING
Jaffe, Bernard. Chemical Calculations. World Book Co.,
Yonkers, 1947. A systematic presentation of the solution of
type problems, with 1000 problems arranged progressively
according to lesson assignments.
Kendall, James. At Home among the Atoms. D. Appleton-
Century Co., New York, 1932. "A Few Figures" explains atomic
weights in a novel way. Tells how atomic weights are found.
USEFUL IDEAS DEVELOPED
1. The atomic weight of an element is a number that rep-
resents the ratio of the weight of 1 of its atoms to the weight
of 1 atom of oxygen. Atomic weights are all relative weights.
2. The molecular weight of a compound is the ratio ol the
weight of 1 molecule of the compound to the atomic weight of
oxygen (16) .
3. A chemical symbol, in addition to representing an ele-
ment, represents one atomic weight of that element.
4. A mole of a substance is its molecular weight expressed
in grams.
USING WHAT YOU HAVE LEARNED
Group A
1. (a) Who issued the first table of atomic weights? (b)
Why was it later decided to use the weight of the oxygen atom
as a standard instead of the weight of the hydrogen atom?
2. Exactly what is meant by saying that oxygen has an at.
wt. of 16?
3. Find the mol. wt. of the compounds that have the fol-
lowing formulas: (a) cupric acetate, Cu (C2H3O2) 2 • H2O;
(b) chloroplatinic acid, H2PtCl2; (c) microcosmic salt,
HNaNH4PO4-4H2O. Refer to Table 4 on page 127.
136 NEW WORLD OF CHEMISTRY
I ...
4. Find the percentage composition of each of the follow-
ing compounds: (a) BaCO3, (b) KMnO4, (c) K4Fe (CN) 6.
Check each result.
5. Determine the percentage composition of each of the
following compounds: (a) BaSO4; (b) KCN; (c) (NH4) 2CO,.
6. Calculate the percentage of H0 in alum,
KA1(S04)2-12H20.
7. Find the percentage of water of crystallization in
Sr (N03) 2 • 5H20.
8. A ton of limestone, CaCOa, was heated in a lime kiln
until all of it was changed to quicklime, CaO. The equation
for this reaction is: CaCO3 -» CaO -+- CO2j. How much
quicklime was formed? To answer this question, first decide
what type ot problem this is: What four steps have you learned
to take in solving such a problem? Solve the problem.
9. Find the weights of (a) H2 and (b) ZnSO4 formed by
the complete reaction of 130 g. of Zn and sufficient H.,SOt.
Group B
10. 320 g. of Fe^Oy, on being reduced, form 224 g. of Fe.
What is the at. wt. of oxygen?
11. 11.95 g. of lead sulfide, PbS, will produce 10.35 g. of
lead. From this fact, and knowing that the at. wt. of S is 32,
calculate the at. wt. of Pb.
12. Of what use to the manufacturing chemist is knowledge
of the percentage composition of a compound? Select from
the chapter on Cu an ore of that metal and illustrate.
13. Suppose a chemist were going to manufacture HC1 from
NaCl and H2SO4. What helpful information could he gain
from the following equation representing the reaction that
would occur? 2NaCl + H2SO4 >Na2SO4 + 2HC1
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Consult a manufacturing chemist or an analytical chemist
and discuss with him how he uses chemical mathematics in
his business or profession. Report your findings to your class.
2. Construct a large graph to represent the percentage of
each element in crystallized washing soda, Na2CO3 • 10H2O.
Use different colors for each element. Show also the per-
centage of H«Q*8JP;this compound.
CHLORINE
AND THE HALOGEN FAMILY
. . . Search for the truth is in one
way hard and another easy, for it is
evident tliat no one can master it
; fully nor miss it wholly. But each
adds a little to our knowledge of Na-
ture and from all the facts assembled
there arises a certain grandeur.
Aristotle
A Swedish druggist discovers chlorine. One of the scientists who
lived and worked at about the same time as Priestley, Cavendish, and
Lavoisier was a Swedish pharmacist who is well known to the world
of chemistry. He not only prepared oxygen earlier than Priestley,
but also, in 1774, discovered chlorine. Scheele is the only great chem-
ist whose whole lifework was accomplished behind the counter or
in the prescription laboratory of one drugstore or another!
When he had discovered the greenish-yellow chlorine gas that
made his nose and throat sting and almost blinded him, he wrote to
a friend: "Oh how happy I am; 1 seldom think of eating, or drinking,
or where I live; I scarcely pay attention to my pharmaceutical busi-
ness. But to watch new phenomena, this is my consuming interest."
At 43, as a result of prolonged exposure to unhealthful conditions in
his crude laboratory, Scheele died — a martyr to the rapidly develop-
ing and expanding science of experimental chemistry.
How chlorine is prepared in the laboratory. Hie method used by
Scheele to prepare chlorine is still the common laboratory method
used today. Two compounds are used — manganese dioxide and
hydrochloric acid. When this mixure is heated, chlorine is liberated
from the hydrochloric acid.
137
1S8
MnO2+HCl
Fig. 21. Laboratory preparation of chlorine. What
property of chlorine makes it wise to pass the excess
gas into water?
Although chlorine gas is fairly soluble in water, it may be col-
lected by water displacement. Since it is heavier than air, it is us-
ually collected by displacing air (see Fig. 21) .
The concentrated hydrochloric acid used in this reaction supplies
the chlorine. In this reaction the manganese dioxide acts as an oxi-
dizing agent, combining with the hydrogen of the acid and liberating
free chlorine. The equation for this reaction is:
[4HjCl + MnJO^i -» 2H2O + MnCl2 + C12
This equation may be considered to represent two reactions. The
first is of the double-replacement type:
MnO2 + 4HC1 -» MnCl4 + 2H2O
and the second, a simple decomposition: *.
MnCl4 -» MnCl2 + C12
Other oxidizing agents such as potassium chlorate or lead diox-
ide, PbO2, may be used instead of manganese dioxide.
The physical properties of chlorine. Chlorine is a greenish-yellow
gas, two and one-half times as heavy as air. It is fairly soluble in
water, forming yellowish chlorine water (one volume of water dis-
solves about two volumes of chlorine gas under normal laboratory
conditions) . It has a penetrating odor and attacks the membranes of
the nose, throat, and lungs. Inhaling ammonia or alcohol vapor
counteracts this irritating action to some extent. Chlorine was the
first gas to be liquefied. It liquefies at about — 34°C. at normal
pressure. Faraday, who was the first to liquefy it, wrote to a friend
in 1823, "I hope to be able to reduce many other gases to the liquid
form." He did.
CHLORINE AND THE HALOGEN FAMILY
139
Chemical properties of chlorine. Chlorine is a typical nonmetal.
It has a valence of one, that is, it combines with monovalent hydro-
gen, atom for atom. Chlorine is very active chemically. It unites
with nearly all other elements, forming compounds called chlorides,
just as oxygen forms oxides. Thus, a bit of sodium reacts brilliantly
when heated with moist chlorine.
Here is a thrilling example of the marvels of chemistry. Sodium, a
silvery, poisonous solid, unites with chlorine, a greenish-yellow,
poisonous gas, producing a white solid, common salt, which is essen-
tial in the diet of both man and all animals.
2Na + C12 -> 2NaCl
Chlorine has a strong attraction for hydrogen. When hydrogen
and chlorine are mixed and exposed to strong light or ignited by a
spark, they combine with explosive violence, forming hydrogen
chloride gas. The equation for this reaction is:
H2+C12->2HC1
If a jet of hydrogen burning in air is thrust into a jar of chlorine,
it will continue to burn, giving off hydrogen chloride as the product
of combustion.
So powerful is the attraction of chlorine for hydrogen that it will
tear hydrogen away from some of its compounds. Thus, when tur-
pentine, a compound of carbon and hydrogen, is poured over a piece
of filter paper which is then thrust into a bottle of chlorine, a flash
of light occurs and a black powder is formed. The black powder is
the free carbon which remains after the hydrogen from the turpen-
tine has combined with the chlorine, forming hydrogen chloride.
Chlorine will combine with water, liberating oxygen.
In this industrial pla
chlorine is being man
factored in mercury cells.
140 NEW WORLD OF CHEMISTRY
How chlorine is prepared for industrial use. The electrolysis of
brine, a water solution of sodium chloride, is the source of most of
the chlorine used today. The electric current liberates free chlorine
at the carbon anode. At the cathode the sodium liberated reacts
immediately with the water, forming free hydrogen and sodium
hydroxide.
2NaCl -> C12 + 2Na
2Na + 2H2O -* 2NaOH + H2
The combined equation, then, is:
2NaCl + 2H2O -» C12 1 + H2 \ + 2NaOH
The chlorine gas is dried by passing it through concentrated sul-
luric acid, and then liquefied. The yellow, liquid chlorine is stored
in steel cylinders, each holding from one to 300 pounds of chlorine
free from water vapor. This process, by which three valuable prod-
ucts— chlorine, hydrogen, and sodium hydroxide — are formed
from two low-cost, plentiful compounds by means of an economical
outlay of electric energy, is described in greater detail on page 212
under the industrial preparation of sodium hydroxide. Some chlo-
rine is also obtained by the electrolysis of molten NaCl (see page
378).
The scientist serves humanity. Soon after the discovery of chlorine,
Berthollet hit upon the idea of using the bleaching action of chlorine
(which Scheele had noticed) industrially. He declined to patent his
process or make any profit from it, but instead turned it over to the
French government. This action of Berthollet is characteristic of
many scientists, who believe that because they were freely helped
toward making their discoveries by the work of those who preceded
them, so they should freely pass on the results of their own labors.
Fig. 22. Chlorine bleaches indirectly. In the presence of sunlight
chlorine reacts with water, liberating nascent oxygen.
chlorine water chlorine wat^r ir]^iWi^)t
oxygen gas-
CHLORINE AND THE HALOGEN FAMILY 141
Patents on chemical processes are sometimes turned over by their
discoverers to foundations, which license them to manufacturers.
These foundations receive fees or royalties for the use of the process.
These funds may be used by the foundations in making further
research possible.
Chlorine bleaches paper and textiles. "Dephlogisticated marine
acid air," as chlorine was called before Davy showed it to be an
element, was a chemical curiosity in 1774. Today, it is an indus-
trial necessity, and more than two million tons of it are produced in
this country annually.
The chief use of chlorine is in the bleaching, or decolorizing, of
textiles, chiefly cotton and linen, and of wood pulp. It cannot be
used for bleaching silk or wool because it destroys their fibers. Chlo-
rine bleaches indirectly, by oxidation. Oxygen, liberated by the ac-
tion of chlorine on water, combines chemically with certain coloring
matters and dyes, which, upon oxidation, become colorless. This re-
action is:
??? +!?• ° ~* 2HC1 + ° t
Chlorine is used as a bleaching agent either in the free condition,
in chlorine water, or in some unstable chlorine compounds, such as
bleaching powder, CaOCl2, and calcium hypochlorite, Ca (CIO) 2.
Laundry bleach, used extensively today, is made by adding liquid
chlorine to a very cold solution of sodium hydroxide. This product
is known as Javelle water and contains sodium hypochlorite, NaCIO,
a salt of hypochlorous acid, HC1O.
2NaOH + C12 -* NaCIO + NaCl + H2O
Sodium hypochlorite decomposes easily, liberating atomic oxygen.
Dilute solutions of sodium hypochlorite, in strengths ranging from
four to six percent, are marketed as "Clorox," "Rose-X," and a num-
ber of other trademarks. Such household bleaches are used very
widely in home laundering and in removing stains. In bleaching with
chlorine, care must be taken not to expose the materials to the action
of the chlorine for too long a time, since continued action will make
the fibers very weak. After the bleaching agent has been used, the
fabrics should be thoroughly rinsed in fresh water to remove all
traces of the bleach. In industrial bleaching, after removal from the
bleaching tank, an antichlor, such as sodium thiosulfate (Na2S2O.<) ,
commonly called hypo, is added to remove the excess chlorine, and
then the material is thoroughly rinsed in running water.
142 NEW WORLD OF CHEMISTRY
Chlorine saves lives. Although it consumes only about ten percent
of normal production, a second very important use of chlorine is in
the purification of water. When chlorine is added to water, atomic
oxygen, liberated from the water by chlorine, reacts with the bacteria
present and kills them by oxidation. Only three pounds of chlorine
are used for each million gallons of water. In the United States more
than 75 percent of all drinking water is treated with chlorine. As a
result, the death rate from typhoid fever, a disease caused by typhoid
bacilli which may be present in drinking water, has been cut down
considerably. This treatment also kills algae and other low plant
and animal life. For field use, in areas where pure water is not avail-
able, explorers, scouts, and others carry tablets of a chlorine-produc-
ing substance such as "halazone" which contains Ca (CIO) 2.
Chlorine is also used as an antiseptic and disinfectant. A substance
that checks the growth of bacteria is called an antiseptic, while a
germicide kills the bacteria outright. A substance that either kills or
checks the growth of bacteria is called a disinfectant. Since much ill-
ness is caused by poisons, or toxins, formed by bacteria, it is impera-
tive that these bacteria be killed or at least prevented from mul-
tiplying.
"Zonite," a trade preparation containing NaCIO, is used as a
general household disinfectant. Because chlorine disinfectants if used
improperly may destroy body tissue, such preparations should be used
with great care. Several widely used insecticides are chlorine com-
pounds (see page 540) .
And chlorine destroys human life, too. Chlorine played a double
role during World War I. While chlorine disinfectants and water
chlorination were saving thousands of lives, free chlorine in the
form of clouds of poisonous gas was choking out many other
lives. Chlorine, and later phosgene, COCL, and mustard gas,
(C1CH2CH2) 2S, caused terrible suffering in World War I, even
though the use of poisonous gases had been "outlawed" by a confer-
ence at The Hague in 1907.
How to test for the presence of a chloride. Free chlorine may be
recognized by its characteristic color and odor, but compounds of
chlorine cannot be identified so easily. Because chlorides are presejat
in so many common substances, a simple test for a chloride is de-
sirable.
All chlorides are soluble in water, with the exception of the
chlorides of silver, mercury, and lead. Because silver chloride is in-
soluble in water, chlorides are recognized by their reaction with a
solution of silver nitrate. When a solution of silver nitrate is added
143
silver nitrate
Fig. 23. Testing an unknown solution for
the presence of a chloride. What will hap-
pen to the white precipitate on exposure
to light if a chloride is present?
to a solution of a chloride, a white insoluble substance, a precipitate,
forms. On exposure to light, the color of this precipitate changes
gradually to purple and then to black.
AgNO3 + NaCl -» NaNO3 + AgCl |
This color change, when silver nitrate is added to an unknown
solution, suggests that probably a chloride is present. But some
substances that are not chlorides form similar compounds. Therefore,
the formation of such a precipitate on the addition of silver nitrate
to a solution is not an entirely reliable test for a chloride.
A chemical test for a substance has one important requirement —
it must be specific, that is, no other substance will react to the test in
the same way as the substance for which the test was designed. For-
tunately, silver chloride is insoluble in dilute nitric acid; other sub-
stances that might at first be mistaken for it are soluble in dilute
nitric acid and disappear at once when this acid is added. The addi-
tion of nitric acid, therefore, is the final step in this test, distinguish-
ing the chloride from other compounds.
An element is known by the company it keeps. The halogens (salt
formers) are a group of elements that resemble one another chemi-
cally, and whose physical properties differ from one another in
regular gradation, as shown in Table 5. Such a group of elements
is called a chemical family. The members of the halogen family are
fluorine, chlorine, bromine, iodine, and astatine, a radioactive ele-
ment. Table 5 shows the relationship among the members of the
halogen family.
Making bromine in the laboratory. Like chlorine, bromine is pre-
pared by the oxidation of its hydrogen compound by manganese diox-
ide. A mixture of potassium bromide, sulfuric acid, and manganese
dioxide is heated in a test tube (see Fig. 24) . The H2SO4 reacts
with the KBr, forming HBr, which is then oxidized by the MnO2.
Although at room temperature bromine is a liquid, it is liberated as
•MMBM^HlHHHi
ELEMENT
ATOMIC
WEIGHT
IAPPROX.]
PHYSICAL PROPERTIES
r
STATE
COLOR
ODOR
SOLUBILITY
IN WATER
^
BOILING
POINT
FLUORINE, F2 1
19
Gas
Pale
yellow
Penetrating
Decomposes
water
-187°C
CHLORINE, C!2I
35.5
Gas
Greenish -
yellow
Irritating
Fairly
soluble
About
-34°C
BROMINE, Br2 1
80
Liquid
Red
Suffocating
Fairly
soluble
About
61°C
IODINE, I2 1
127
Solid
Purplish -
black
Resembles
chlorine
Least soluble
of halogens
Above
200°C
a brownish vapor at the temperature of the experiment. As this
brownish vapor passes into the water, some of the bromine dissolves;
the rest collects as a layer of bromine under the water. This method
of collecting pure bromine by distillation is relatively safe.
2KBr + 2H2SO4 + MnO2 -> K2SO4 + MnSO4 + 2H2O + Br2
Great care should be taken in working with bromine, because it is
poisonous and attacks the skin, causing severe burns. Particular care
should be taken to protect the eyes from bromine vapor.
Taking bromine from the sea. Most bromine is extracted from the
minute percentage (0.0065%) of bromides present in sea water.
Free chlorine replaces the bromine of the bromides. Some of our
bromine is also obtained from the bromides found in salt wells and
salt lakes. The principal chemical reaction is:
MgBr2 + C12 -» MgCl2 + Br2
Bromine helps engine efficiency. Much of the bromine produced
in the United States is used in the manufacture of "Ethyl fluid,"
an anti-knock mixture composed of ethylene dibromide and tetra-
ethyl lead (TEL) . Large quantities of bromine are also used in
making silver bromide, the light-sensitive chemical that forms the
most important part of the coating of photographic films. Bromine
and bromine compounds are used also in making tear-gas bombs.
Bromine is used in appreciable quantities as an oxidizing agent in
the manufacture of certain dyes and drugs.
Potassium bromide or sodium bromide acts as a depressant on the
central nervous system. Their action is followed by drowsiness and
even sleep. Such chemicals are called sedatives. They are used in the
144
CHEMICAL PROPERTIES
DISCOVERY
CHIEF
ORE
CaF2
NaCI
MgBr2
Nal03
Mgl2
r
VALENCE
ACTION
WITH
HYDROGEN
ACTION
WITH
METALS
ACTION
WITH
WATER
ORDER
OF
ACIVITY
One
Forms HF
Forms
fluorides
H20 + F2
-*2HF + 0
Most
active
Moissan
!mwa-saN'),
1886
One
Forms HCI
Forms
chlorides
H20 + CI2
-*2HCI + 0
Next
active
Scheele,
1774
One
Forms HBr
Forms
bromides
Br2-t-H2O
-* 2HBr + O
Less than
chlorine
Balard
(ba-lar'l,
1826
One
Forms HI
Forms
iodides
I2 + H20
-» 2HI + 0
Least
active
Courtois
(koor-twa'),
1812
TABLE 5
THE HALOGEN FAMILY
treatment of insomnia and asthma and are frequently found in head-
ache and sleeping powders. Heavy, continuous doses of bromides may
have harmful effects on the body. Bromides should be used only on
the advice of a physician.
Methyl bromide, CH;5Br, is used widely in commercial and indus-
trial fumigation to kill insects and other low forms of plant and ani-
mal life. It is used for this purpose in boxcars, warehouses, and food
processing and packaging plants.
How iodine is prepared in the laboratory. Like chlorine and
bromine, iodine is prepared by the oxidation of its hydrogen com-
pound by manganese dioxide. A mixture of potassium iodide, sulfuric
acid, and manganese dioxide is heated in a test tube as shown in
Fig. 25.
SKI + 2H2S04 + Mn02 -» K2SO4 + MnSO4 + 2H2O -f I2
The violet vapor of iodine, which is produced at the temperature
of the reaction, condenses, forming purplish-black crystals on the in-
side of the test tube. This process of collecting iodine is called sub-
Fig. 24. Preparation of bromine in
the laboratory. Why is bromine not
collected in the same way as chlo-
rine?
bromine
water
bromine
ELEMENT
ATOMIC
WEIGHT
(APPROX.)
PHYSICAL PROPERTIES
••
CHEMICAL PROPERTIES
^•H
DISCOVERY
•i
CHIEF
ORE
STATE
COLOR
ODOR
SOLUBILITY
IN WATER
^\
BOILING
POINT
r
VALENCE
ACTION
WITH
HYDROGEN
ACTION
WITH
METALS
ACTION
WITH
WATER
ORDER
OF
ACTVITY
FLUORINE, F2 1
19
Gas
Pale
yellow
Penetrating
Decomposes
water
- 187°C
One
Forms HF
Forms
fluorides
H20 + F2
-* 2HF + O
Most
active
Moissan
(mwa-saN'),
1886
CaF2
CHLORINE, CI2I
35.5
Gas
Greenish-
yellow
Irritating
Fairly
soluble
About
-34°C
One
Forms HCI
Forms
chlorides
H20 + CI2
-*• 2HCI + O
Next
active
Scheele,
1774
NaCI
BROMINE, Br2 1
80
Liquid
Red
Suffocating
Fairly
soluble
About
61°C
One
Forms HBr
Forms
bromides
Br2 + H2O
-* 2HBr + O
Less than
chlorine
Balard
(ba-lar'),
1826
MgBr2
IODINE, I2 f
127
Solid
Purplish -
black
Resembles
chlorine
Least soluble
of halogens
Above
200° C
One
Forms HI
Forms
iodides
I2 + H2O
-* 2HI + 0
Least
active
Courtois
(koor-twa'J,
1812
NalO3
Mgl2
a brownish vapor at the temperature of the experiment. As this
brownish vapor passes into the water, some of the bromine dissolves;
the rest collects as a layer of bromine under the water. This method
of collecting pure bromine by distillation is relatively safe.
2KBr + 2H2SO4 + MnO2 -» K2SO4 + MnSO4 + 2H2O + Br2
Great care should be taken in working with bromine, because it is
poisonous and attacks the skin, causing severe burns. Particular care
should be taken to protect the eyes from bromine vapor.
Taking bromine from the sea. Most bromine is extracted from the
minute percentage (0.0065%) of bromides present in sea water.
Free chlorine replaces the bromine of the bromides. Some of our
bromine is also obtained from the bromides found in salt wells and
salt lakes. The principal chemical reaction is:
MgBr2 + C12 — > MgCl2 + Br2
Bromine helps engine efficiency. Much of the bromine produced
in the United States is used in the manufacture of "Ethyl fluid,"
an anti-knock mixture composed of ethylene dibromide and tetra-
ethyl lead (TEL) . Large quantities of bromine are also used in
making silver bromide, the light-sensitive chemical that forms the
most important part of the coating of photographic films. Bromine
and bromine compounds are used also in making tear-gas bombs.
Bromine is used in appreciable quantities as an oxidizing agent in
the manufacture of certain dyes and drugs.
Potassium bromide or sodium bromide acts as a depressant on the
central nervous system. Their action is followed by drowsiness and
even sleep. Such chemicals are called sedatives. They are used in the
144
TABLE 5
THE HALOGEN FAMILY
treatment of insomnia and asthma and are frequently found in head-
ache and sleeping powders. Heavy, continuous doses of bromides may
have harmful effects on the body. Bromides should be used only on
the advice of a physician.
Methyl bromide, CH,,Br, is used widely in commercial and indus-
trial fumigation to kill insects and other low forms of plant and ani-
mal life. It is used for this purpose in boxcars, warehouses, and food
processing and packaging plants.
How iodine is prepared in the laboratory. Like chlorine and
bromine, iodine is prepared by the oxidation of its hydrogen com-
pound by manganese dioxide. A mixture of potassium iodide, sulfuric
acid, and manganese dioxide is heated in a test tube as shown in
Fig. 25.
2KI + 2H2SO4 + MnO2 -> K2SO4 + MnSO4 + 2H2O + I2
The violet vapor of iodine, which is produced at the temperature
of the reaction, condenses, forming purplish-black crystals on the in-
side of the test tube. This process of collecting iodine is called sub-
KBr + MnO2
H2SO4
Fig. 24. Preparation of bromine in
the laboratory. Why is bromine not
collected in the same way as chlo-
rine?
bromine
water
iHk" bromine
Dow Chemical Company
Methyl bromide used as a fumigant destroys all grain insects.
limation. A substance is said to sublime when it passes directly from
the solid state to the gaseous state and then condenses back to the
solid state without passing through the liquid state. Camphor1 is an-
other substance that sublimes.
The industrial preparation of iodine. About 90% of our iodine
is obtained from brine that comes up with the oil in California oil
fields. This brine contains Nal and MgL. Chlorine is passed through
the brine and replaces the iodine. The iodine is recovered by adsorp-
tion on activated carbon. (Adsorption is the clinging of molecules of
one substance to the surface of another — see page 327.) The princi-
pal reaction is:
MgI2 + C12 -> MgCl2 + I2
The rest of our iodine comes from NaIO;5, found as an impurity
in Chile saltpeter, NaNO,.
Iodine, too, saves lives. The chief use of iodine is in the prepara-
tion of tincture of iodine, a two percent solution of iodine and po-
tassium iodide in ethyl alcohol, which is an excellent antiseptic. As
silver iodide, iodine is used to some extent in photography together
with silver bromide. It is used also in the manufacture of iodoform,
CHI:1, a yellow powder used as an antiseptic, and in the manufacture
of "Aristol," an improvement over iodoform. Iodine is also used in
the production of certain dyes and methyl iodide, CHJ.
146
CHLORINE AND THE HALOGEN FAMILY
147
Is iodine necessary to health? Iodine is an important constituent
of the human body. There is a definite relation between the presence
of iodine in the thyroid gland and the prevalence of certain disorders.
The thyroid gland, located in the neck, secretes a compound called
thyroxin, containing about 65 percent iodine, which helps to regulate
the rate of oxidation in the body.
When the thyroid gland receives too little iodine, goiter, an en-
largement of the thyroid, results, caused apparently by the attempt of
the gland to increase its size in order to produce more thyroxin. To
offset this deficiency, iodides may be added to drinking water or
about 0.02 percent of sodium iodide added to so-called iodized salt.
Extreme underactivity ot the thyroid gland in newborn babies and
young children may result in cretinism (kre'tm-izm) — misshapen
dwarfishness, low mentality, sluggishness, dullness, slow heart action.
Synthetic thyroxin is used in the treatment of this thyroid disorder.
Overactivity of the thyroid gland often produces the opposite effect —
the thin, nervous, highly energetic person, whose movements are
quick, and whose heart action is rapid. See page 34 for a discussion
of basal metabolism tests in diagnosis of thyroid disturbances.
Iodine is also necessary to other forms of animal life. Large quan-
tities of iodides are added to commercial feeds for chickens, cattle,
dogs, cats, and other animals, and to fertilizers for forage crops.
Replacement power of the halogens. If free chlorine is added to a
solution of a bromide or an iodide, free bromine or free iodine is
liberated. Free chlorine replaces the two less active halogens.
2KBr + C12 -> 2KC1 + Br2
2NaI + C12 -> 2NaCl + I2
The addition of free bromine to a solution of an iodide releases
free iodine.
2KI + Br2 -* 2KBr + I2
test tube containing I
' H2SO4
<^:2
Fig. 25. Laboratory preparation of
iodine.
iodine crystals
148 NEW WORLD OF CHEMISTRY
However, the addition of iodine to a solution of either a bromide or
a chloride has no effect, for the less active halogen will not replace
the more active halogen from its compound. As mentioned earlier,
the commercial preparation of bromine depends on the replacement
power of chlorine.
How we test for the presence of bromides and iodides. Many uses
are made of the replacement power of the halogens. The tests for
bromides and iodides are based on it. Chlorine water is added to a
solution of the unknown salt, and a few drops of carbon disulfide,
CS2, which is not soluble in water, are also added. If, after shaking
the mixture, the carbon disulfide settles out as a distinct layer with
a brownish-red coloration, then the original salt was a bromide, the
free bromine liberated coloring the carbon disulfide. If the carbon
disulfide acquires a purple coloration, the original salt was an iodide.
Carbon disulfide is used because free bromine and iodine are much
more soluble in this liquid than in water. Hence, most of the liber-
ated bromine and iodine dissolve in the carbon disulfide, thus color-
ing it much more than they would color water. Carbon disulfide, be-
cause it is a better solvent, will extract any bromine or iodine from
the water solution. This process of separation, frequently used in
industry, is called extraction by partition.
Fluorine, the most active of all the elements. Fluorine was not
isolated until 1886. Because of its extreme chemical activity, which
causes it to unite violently with metals, glass, porcelain, and water,
its separation as a pure element was a very difficult undertaking.
Finally, Henri Moissan succeeded by liquefying pure hydrogen
fluoride, adding some potassium fluoride, and at a temperature of
— 23°C. passing an electric current through the mixture. Fluorine
was liberated at the anode.
The anode used industrially today is made of graphite, which is
not attacked by this pale yellow gas. Fluorine is employed in making
uranium hexafluoride used in atomic energy plants where many of
the lubricants are chemically inert fluorocarbon compounds. The
plastic, "Teflon," is another fluorine compound. The new rat and
ground squirrel poison, 1080, is a fluorine compound, and sodium
fluoride is used in some insecticides.
Fluorine prevents tooth decay. The amount of tooth decay, or
dental caries, has been found to vary directly with the quantities of
fluorides present in the- local water supply. Too much fluoride pro-
duces very hard but mottled teeth, a condition in which the enamel
becomes discolored badly. When too little fluoride is present, there
is much tooth decay. Water fluoridation (about one part NaF per
CHLORINE AND THE HALOGEN FAMILY
149
million parts water) is now widely practiced to protect children up
to about the age of 12 while tooth enamel is being formed.
Fluorine refrigerants. A refrigerant is a substance used to absorb
heat by changing from a liquid to a gas. In refrigeration, the material
from which heat is absorbed is cooled. Almost as long as mechanical
refrigerators have existed, their manufacturers have searched for bet-
ter refrigerants. Something was wrong with nearly all the original
refrigerants. They were either toxic, flammable, corrosive, or pos-
sessed disagreeable odors. And then a family of compounds was de-
veloped, and introduced in 1931 under the trademark "Freon."
These compounds, produced by the halogenation of simple com-
pounds of carbon and hydrogen, are far superior to sulfur dioxide,
ammonia, ethyl chloride, and methyl chloride as refrigerants. All are
practically odorless, nontoxic, nonflammable, and noncorrosive.
The first of the "Freons" to be produced was dichlorodifluoro-
methane, CC1,F,, and a later one was dichloroteArafluoroethane,
C2C12F4. There are several others, each with slightly different proper-
ties that make it particularly well adapted to a special use. Because of
their extreme volatility and the speed with which they penetrate
every nook and cranny of a confined space, the "Freons" are used as
propellants in dispersing insecticides. Aerosol bombs used for killing-
household pests usually contain an insecticide and a liquefied "Freon"
gas under pressure. When the pressure is released, the expanding gas
quickly distributes the insecticide throughout a room or closet.
This electronic device, called a "sniffer/1 is used to detect
breaks in telephone cables through which moisture might
enter. The cable to be tested is filled with a "Freon1' gas.
Then the sniffer is pulled along the cable from the ground.
Escaping gas activates a "FreorT'-sensitive detector, ringing
a bell.
Bdl Telephone laboratories
|P^'^-^:^^^^
150 NEW WORLD OF CHEMISTRY
YOU WILL ENJOY READING
Consumer Reports (Consumers Union Reports) , April, May,
June, 1939. Consumers Union of U.S., New York. Excellent
reports on various antiseptics.
Harrow, Benjamin. Eminent Chemists of Our Time (2nd
ed.) . D. Van Nostrand Co., New York, 1927. Read the life of
Moissan and his isolation of fluorine.
Sanders, Gardiner, and Wood. "Chlorine and Caustic Manu-
facture/' Industrial and Engineering Chemistry, September,
1953, pp. 1824-1835. Includes history, production figures,
photos, and diagrams showing how diaphragm and mercury
cells work.
USEFUL IDEAS DEVELOPED
1. The halogens are a group of elements that resemble one
another chemically, and whose physical properties differ in
regular gradation. The halogen group is one of several such
groups of elements.
2. A substance that passes directly from the solid to the
gaseous state and directly from the gaseous to the solid state
is said to sublime.
3. The ability of one element to replace another in that ele-
ment's compounds is widely used in chemical reactions. The
halogens may be listed according to their replacement power.
4. The separation of a substance from a solution containing
that substance by means of a better solvent is known as extrac-
tion by partition.
5. Like Berthollet, many scientists believe they should freely
pass on to others the results of their own labors, and so help
all humanity.
USING WHAT YOU HAVE LEARNED
Group A
1. Who discovered C12, and in what year was it isolated?
2. Make a labeled diagram showing the laboratory prepara-
tion of C12.
3. What is the function of MnO2 in the laboratory prep-
aration of C12?
4. What other substance might be used instead of MnO2 in
the preparation of C12?
CHLORINE AND THE HALOGEN FAMILY 151
•<5. (a) .Give two reasons for collecting C12 by the displace-
ment of air. (b) Why should any excess of C12 be passed into
water?
6. Write a balanced equation representing the laboratory
preparation of Cla.
t . . .
7. State four physical properties of C12.
8. Faraday liquefied C12 in 1823. It was the first gas to be
liquefied. Suggest a reason.
9. If a brightly burning paraffin taper is inserted in a bottle
of C12, a heavy black smoke is given off. Explain.
10. By what process does C12 bleach? Explain.
11. (a) Why is it necessary to rinse materials that have
been bleached with C12? (b) Why cannot C12 be used to bleach
silk and wool?
12. Because of what chemical property are bleaching pow-
der, CaOCl2, and Clorox, containing NaOCl, able to bleach?
13. Find the percentage of C12 in bleaching powder.
14. (a) Distinguish between the terms antiseptic and germi-
cide, (b) What term embraces both?
15. Make a list of all the uses of C12 that you know.
16. Chloride of lime, or bleaching powder, has the formula
CaOCl2. It is made from Ca (OH) 2 and C12. Write the equa-
tion for its preparation.
17. Give a brief account of the preparation of some laundry
bleach.
18. Chlorine played a double role in wartime. Explain.
19. How much CI2 can be prepared by the action of 348 g.
of MnO2 on sufficient concentrated HC1?
. . t . . .
" 1
20. (a) Describe fully the test for a chloride, (b) Write the
equation for the reaction that takes place.
21. Illustrate the statement, "An element is known by the
company it keeps."
22. (a) Compare the physical properties of the members of
the halogen family, (b) Compare their chemical properties.
23. In what ways do the halogens (a) resemble one an-
other, and (b) differ from one another?
24. List the halogens in the order of their chemical activity.
152 NEW WORLD OF CHEMISTRY
25. Illustrate what is meant by the replacement power of
the halogens.
26. Does the following equation represent an actual chem-
ical reaction? Explain. 2KC1 + Br2 >2KBr + C12
27. Describe two uses of the replacement power of the
halogens.
28. Make a labeled diagram of the laboratory preparation
of Br2.
29. (a) What is a sedative? (b) State three other uses of
2'
30. Describe the commercial preparation of Br2.
Brr
31. Using a labeled diagram, describe the laboratory prep-
aration of I2.
32. (a) What is sublimation? (b) What is tincture of
iodine?
33. How is iodine obtained for industrial use?
34. How is the most active of all the chemical elements
prepared?
35. (a) What is one cause of mottled teeth? (b) What is
fluoridation?
36. What is the relationship of F2 to dental caries?
37. (a) What compounds of F2 are superior refrigerants?
(b) Why?
Group B
38. Compare C12 and O2 with respect to: (a) chemical
activity, (b) behavior with H2, (c) valence. Explain each
answer fully.
39. (a) What is the relationship between lack of iodine
and goiter? (b) Why is goiter not so prevalent in New York
City as it is in some other parts of the United States?
40. A bottle of tincture of iodine was found, after long
use, to contain only a dark solid, (a) Would it be safe to use
it after adding pure ethyl alcohol? (b) Explain your answer.
41. If you had some Nal, how would you prepare a solution
of tincture of iodine from it?
42. I2 is produced from California oil-well brines and from
other brines in Michigan. Can you suggest a way in which
this I9 is extracted?
CHLORINE AND THE HALOGEN FAMILY 153
43. How much fluorine would be needed to make one ton
of dichlorotetrafluorethane, C2C12F4?
44. (a) What factors must a manufacturer consider when
he chooses raw materials for use in manufacturing a substance
on a large scale? (b) Show how these factors apply in the man-
ufacture of C12.
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Make a study of advertisements of antiseptics and dis-
infectants in magazines and newspapers. What appeals are
made in these advertisements to induce you to buy a particular
brand? Are the appeals chiefly scientific, pseudo-scientific, fear-
provoking, or do they appeal chiefly to your pride, sense of
superiority, your desire for social approval, and so forth? Il-
lustrate your report with actual advertising copy.
2. Ink eradicators frequently contain two solutions. No. 1
contains a solution of a very weak acid and No. 2 contains a
solution of sodium hypochlorite. Prepare such an ink eradi-
cator using vinegar or citric acid for No. 1 and Javelle water
or Clorox for No. 2. Demonstrate its use before your class.
3. Victor Meyer, an eminent German chemist, prepared a
compound in 1886 which is now known as mustard gas. Meyer
was the first chemist to prepare this chemical during his the-
oretical investigations. Thirty years later chemists, cooperating
with the Germany military machine, made this compound
available for use as a poison gas. Should society take a hand in
suppressing such discoveries which might be used against man-
kind? Write a brief paper either in favor of this point of view
or against it, or arrange in class for a discussion or debate on
this topic.
4. Study the quotation at the beginning of this chapter.
Prepare a brief report to the class on the meaning of this
quotation. Illustrate your report with an example from the
history of science.
ELECTRONS
AND OTHER PARTICLES
. . . Our experiences and observa-
tions alone never lead to finalities.
Theory, however, creates reliable
roads over which we may pursue our
journeys through the world of ob-
servation. Anton Reiser, 1930
The electron theory gives us a clearer picture of matter and its
changes. Formulating accurate theories takes remarkably clear in-
sight, courage, and creative imagination. The theories and principles
of science are among the most truly creative products of the mind of
man. The atomic theory of Dalton is one of the great theories upon
which modern chemistry is built. It shows us that atoms do not com-
bine in a haphazard, irregular manner but form molecules in ac-
cordance with unvarying natural laws. However, it gives us no idea
why this is so.
A more recently developed theory, which is called the electron the-
ory, supplements the atomic theory and gives this explanation. It pro-
vides answers also to such questions as why the extreme chemical
activity of fluorine, the comparative inactivity of nitrogen, and the
inertness of argon; why elements and radicals possess the valences
that they have, and many other questions. The electron theory is the
fruit of many scientists who worked in many countries. It made possi-
ble the atomic age.
Static electricity. About 2600 years ago, the Greeks discovered that
when amber is rubbed with cloth, it becomes capable of attracting
tiny bits of straw or dry leaves. Through the centuries, men have
154
pithball
repelled^
rubber rod
Fig. 26. Demonstration of static electricity.
glass rod
discovered that other materials may be given this property. Glass,
when rubbed with silk, or hard rubber, when rubbed with fur, will
also attract light objects. The force which causes this attraction was
named electricity from the Greek word for amber. Today, we refer
to the electricity caused by friction or rubbing as static electricity.
Benjamin Franklin attempts to explain negative and positive
electricity. Suspend a pith ball by a silk thread. Touch it with a hard-
rubber rod which has been rubbed with fur. As soon as contact is
made, the pith ball will be driven away, or repelled, from the rod.
Then bring near the pith ball a glass rod which has been rubbed
with silk. The ball will be attracted toward the glass rod.
This simple experiment demonstrates that the glass and the hard-
rubber rods were charged with opposite kinds of electricity. When
the neutral pith ball was touched by the hard-rubber rod, it became
charged with the same kind of electricity as the rod. It was then
pushed away from the rod, proving that objects with the same electric
charges repel one another. However, the ball was drawn toward the
glass rod, proving that objects with opposite electric charges attract
one another.
In 1747, Benjamin Franklin, one of the most versatile men Amer-
ica has ever produced, received a static electricity machine Irom a
friend in England. Franklin, son of a soap-maker who had fled from
England because of religious persecution, was then 41 years old. Be-
cause of a very successful business career, he was rich enough to re-
tire from business and devote himself to scientific experimentation.
He had already organized the first scientific society in the New
World, an organization which later became the American Philosophi-
cal Society.
Franklin performed many experiments with his static electricity
machine, and that same year he announced his own views on the
nature of electricity. He wrote: "The electric fire (electricity) is not
created by friction, but collected, being really an element diffused
among matter. The electrical matter consists of particles extremely
subtile. . . . Hence have arisen some new terms among us: we say
B is electrised positively; A, negatively. Or rather B is electrised plus;
155
156 NEW WORLP OF CHEMISTRY
A, minus/' Franklin was the first person to use these present-day
terms in referring to electricity. He later induced his good friend
Priestley to write a history of electricity, and thus, in part, directed
Priestley's scientific career.
Franklin's electric theory was not altogether correct, for he be-
lieved that if a body has too much electricity it is charged positively
(-f ) ; if it has not enough, it is charged negatively (— ); and if it has
just enough, it is neutral. Even though his ideas were not altogether
correct, his reasoning and his terminology for electricity were more
modern than those of any other eighteenth-century scientist. So
great was the creative imagination of Franklin that he came very
close to arriving at the modern concept of the electric nature of mat-
ter, a point of view reached only after some 150 years of further ex-
perimentation.
Today we refer to the charge which was produced on the hard-
rubber rod mentioned above as negative (— ) , and that produced on
the glass rod as positive (+) .
The electron is discovered. Almost a century and a half after
Franklin, William Crookes, an Englishman, studied the effect of
passing a current of high-voltage electricity through a glass tube from
which nearly all the air had been pumped. He noticed that a beam of
light issued from the negative plate, or cathode, of the tube. When a
magnet was brought near the tube, the beam would bend. Since ordi-
nary light is not affected by a magnet, the beam showed a property of
matter rather than of light.
Another Englishman, J. J. Thomson, undertook to explain the
strange behavior of these cathode rays. In 1897, after 20 years of bril-
liant research, he announced his results. He said that cathode rays
are composed of particles of negative electricity, torn away from the
atoms of the air in the tube. To these particles, Thomson gave the
name electrons. The cathode ray was bent because the negative elec-
trons would be attracted by the positive pole and repelled by the
negative pole of a magnet.
Fig. 27. Crookes' tube. Notice how the positive pole
of the magnet deflects the cathode ray.
stream of electrons
Crookes' tube
Joseph John Thomson was
born at Manchester in 1856.
He followed Rayleigh as head
of the Cavendish Laboratory
of Experimental Physics at
Cambridge University.
Brumt font tt c,
The discovery of the electron also explained the phenomenon of
static electricity. When one object is rubbed with another, electrons
are transferred. Thus, when glass is rubbed with silk, the glass loses
electrons to the silk and is left with a positive charge; amber, when
rubbed with fur, takes electrons from the fur and becomes negatively
charged.
Thomson's discovery of the electron completely upset the theory
that the atom is the smallest unit of matter, since electrons were
found in all atoms. It also proved that the idea of an indivisible atom
was inaccurate.
Subsequently, Robert A. Millikan, an eminent American scientist,
succeeded in computing the mass of a single electron. He found it
to be about y^Vr °* ^ie mass °* one hydrogen atom. (Mass is the
amount of matter that a substance contains. It does not vary, as does
lirown Brothers
Robert Andrews Millikan (1868-
1 953) won the Nobel physics prize
in 1923 for his work in isolating
and weighing the electron.
157
158
gas
electron (outside the nucleus)
proton (in the nucleus)
Fig. 28. The hydrogen atom
weight, with the gravitational pull of the earth. However, weight is
dependent upon the amount of matter a substance contains.)
The proton is discovered. After Thomson discovered the electron,
one of his students, Ernest Rutherford, began to ponder over the na-
ture of the rest of the atom. An atom itself is electrically neutral.
Surely, he thought, in the electrically neutral atom there must be
some positive electricity to counterbalance the negative electron.
After much research, which began in 1911, Rutherford determined
that all atoms contain one or more particles of positive electricity
which he named the proton. The proton, which is a positively
charged atom of hydrogen, is 1837 times as heavy as the electron.
The structure of the atom according to the electron theory. Atoms
of the various elements contain different numbers of electrons and
protons. In general, however, the arrangement of these particles fol-
lows a similar pattern in all elements. Rutherford gave us the first
picture of the structure of the atom. It resembles our own solar sys-
tem with its sun and revolving planets. The "sun" of the atom is
called the nucleus. It is composed partly of protons around which,
at a relatively great distance, revolve planetary electrons. A diagram
of the simplest of the atoms, that of hydrogen, is shown above.
It was Rutherford who found that electrons and protons were not
evenly distributed in the atom but that the heavy protons were all
located in the center. He shot helium atoms (alpha particles)
through a cluster of nitrogen atoms and photographed the results
by means of a cloud chamber and special camera devised by
C. T. R. Wilson, another English scientist. He found that only in
an extremely few cases the path or fog track of the helium bullet
was not straight, but was thrown sharply off its course. On the basis
of the volume of the nitrogen atom and the ratio of straight fog
tracks to bent ones, he calculated that the protons must all be con-
Fig. 29. (left) Apparatus for producing the cloud effect.
The upper chamber is filled with nitrogen, (right) The
fog track of an alpha particle.
.water vapor condenses
cloud
effect
results
track of particle
suddenly released
bulb
containing
water. . .
is compressed .
ELECTRONS AND OTHER PARTICLES 159
centrated in the tiny nucleus of the atom. The diameter of the nu-
cleus is about 10o*ooo that of the whole atom. Helium bullets pass-
ing through the rest of the atom met no solid, positively charged
protons.
Size of the atom. Most of us are baffled in attempting to imagine
the size of the particles within the atom and the distances between
them. The diameter of a hydrogen atom is about 40,000,000 inch.
If an atom were magnified about 30 trillion times, its diameter would
be about 10 miles. At the center would be the nucleus, about the size
of a tennis ball. The electrons, each about the size of a hazelnut,
would revolve about the nucleus in orbits in somewhat the same way
as the planets of our solar system revolve about the sun.
This description is actually an oversimplification of the structure
of the atom. However, it tells us an important fact. Atoms, which
compose every element and compound, are largely empty space!
The neutron: A notable scientific prediction. All of the protons
of an atom are located in the nucleus. However, not all of the elec-
trons of an atom are planetary; some electrons, too, are found within
the nucleus. But how can negatively charged electrons and positively
charged protons exist side by side in the nucleus? To explain this,
Dr. W. D. Harkins predicted the existence of a new particle. An
electron within the nucleus, he said, does not exist as a separate par-
ticle, but is combined with a proton forming an electrically neutral
particle which he named the neutron. Since the weight of the elec-
tron is extremely slight, it may be disregarded in figuring the weight
of the neutron. The neutron has been found to have about the same
mass as the proton.
In 1920, the existence of the neutron was theoretically established
by W. D. Harkins of the University of Chicago. Twelve years later,
it was actually discovered by James Chadwick, working in Ruther-
ford's laboratory. The ultimate verification of Harkins' work by
Chadwick demonstrates the value of pure theory and creative imag-
ination in science. For his part in the discovery, Chadwick was
knighted in 1944.
The electron and proton are so close together in the neutron that
the volume of the neutron is millions of times smaller than that of
any atom. It has, therefore, an extremely high density. This fact has
been used to explain the extremely high densities of certain stars.
Since the neutron is electrically neutral, electric forces do not repel
it. Therefore, the neutron has great penetrating powers.
We may say that, in general, all matter is composed of three kinds
of fundamental electric particles: electrons, protons, and \neutrons.
160 NEW WORLD OF CHEMISTRY
Scientists have discovered other electric particles within the atom,
as shown in Table 8, but the three mentioned here are considered
the most important.
How the structure of an atom may be represented graphically.
The atom of each element contains a particular number of electrons,
protons, and neutrons different from the number in the atom of
every other element. Two American scientists, Lewis and Langmuir,
developed a theory of the arrangement of planetary electrons which
explains why each element has different chemical properties.
According to this theory, the electrons outside the nucleus (plane-
tary electrons) arrange themselves in successive rings, or shells. The
first ring is complete when it contains 2 electrons; the second ring
is complete when it contains 8 electrons; the third ring is complete
when it contains 18 electrons; the fourth, when it contains either 18
or 32 electrons. However, the outermost ring never contains more
than eight electrons. According to Lewis and Langmuir, an atom
with 30 planetary electrons will have its first ring complete with
2 electrons, second ring complete with 8 electrons, third ring com-
plete with 18 electrons, and fourth ring incomplete with 2 electrons.
What is the electron structure of an atom with 48 planetary electrons?
The number of protons in the nucleus of the atom is equal to the
number of planetary electrons. This equality keeps the atom elec-
trically neutral. Thus, an atom with 30 planetary electrons will have
30 protons in its nucleus.
Of all the elements, hydrogen has the simplest atom. Its nucleus
consists of one proton. Revolving about this nucleus is one planetary
electron. The nucleus of the helium atom contains two protons (2+)
and two neutrons (2n) ; two planetary electrons (2—) revolve about
it. Fig. 30 shows how we may graphically represent the helium atom
and also the chlorine atom (atomic number 17) .
The periodic table of Mendeleyeff. In 1869, Mendeleyeff (men'-
de-la'ef) , a Russian chemist, published a table of the elements ar-
ranged in order according to their increasing atomic weights. He
noticed, that when arranged in this manner, the elements fell into
Fig. 30. Diagrams of the helium and chlorine atoms.
Helium Chlorine
or
GROUP ^
SERIES
ZERO
1
II
III
IV
V
VI
VII
2
He
Li
Be
B
C
N
0
F
3
Ne
Na
Mg
Al
Si
P
S
Cl
4
A
K
Co
Sc
Ti
V
Se
Br
eight distinct groups. Within each group, the elements have similar
physical and chemical properties.
Let us examine part of this table carefully. Note that hydrogen is
omitted, and that the table begins with helium, the element with the
next heaviest atomic weight. Lithium (Li) , with an atomic weight
of 6.940 follows, and so on through fluorine (19.0) . These eight dis-
similar elements comprise one series or period.
The element following fluorine in order of atomic weight is neon
(Ne) . It has chemical properties similar to those of helium (He) and
falls directly below it in the table. Directly below helium is argon
(A) with similar properties. Thus helium, neon, argon, and certain
succeeding elements comprise Group Zero. This group of elements
is known as the family of inert gases (see Chapter 7) . The elements
in Group VII are known as the halogen family (see Chapter 10) .
The other groups are also made up of closely related elements.
Moseley discovers the law of atomic numbers. Prior to 1912, the
numerical position of an element in a table of atomic weight was
called the atomic number of the element. It occurred to Rutherford
that this number might also represent the number of protons in the
nucleus of the atom. One of his students, Henry G. J. Moseley, un-
dertook to find out whether Rutherford's idea was valid.
Moseley's experiments bore out Rutherford's theory. The atom
of each element was found to contain a number of protons in its nu-
cleus corresponding to the element's numerical position in the peri-
odic table of atomic weights. The hydrogen atom, which appears first
in the periodic table, has only one proton in the nucleus of its atom;
uranium, which appears in the ninety-second position in the periodic
table, has 92 protons in the nucleus of each atom. Moseley showed
that the atomic number of any element is equal to the number of free
protons in the nucleus of its atom. This is known as the law of atomic
numbers. Since the number of free protons in the nucleus is equal
to the number of electrons around the nucleus, the atomic number
of an element is also equal to the number of planetary electrons.
For the first 17 elements in the table of atomic numbers, it is help-
ful to remember that the atomic number is equal to half the atomic
161
162
NEW WORLD OF CHEMISTRY
ELECTRONS AND OTHER PARTICLES
163
weight (disregarding fractions). Thus the atomic number of chlo-
rine (atomic weight 35.457) is 17.
The periodic table of atomic numbers. Mendeleyeff's periodic
table based upon atomic weights served science for 50 years. In 1912,
however, it was displaced by the new periodic table developed by
Moseley from his law of atomic numbers.
Moseley's table is more fundamental than MendeleyefFs and easily
accounts for some of the discrepancies in the latter. For example, the
element argon has an atomic weight of 39.944 and the element potas-
sium an atomic weight of 39.10. Argon should, therefore, follow
potassium in the table based on atomic weights. But the properties
of argon put it in the group of inert elements, preceding potassium.
Moseley's researches, which showed that the atomic number of argon
is 18 and that of potassium is 19, eliminated this problem. When
Moseley developed his table, all of the elements had not been discov-
ered. Therefore the atomic numbers did not run entirely consecu-
tively from one to 92. Since then the missing elements have been dis-
covered, as you may see from the modern periodic table below.
A new definition of atomic weight. As we have learned, the elec-
tron, for most purposes, may be considered weightless, and the proton
and neutron may be considered equal in mass. Thus, the most abun-
dant hydrogen atom, which contains a single proton and no neutrons,
may be considered equal to one proton in weight; the helium atom,
which contains two protons and two neutrons, weighs four times as
much as the hydrogen atom. We may say that the atomic iveight of
hydrogen is one and the atomic weight of helium is four. Thus, ac-
cording to the electron theory, the atomic weight for each element
may be defined as the sum of the protons and neutrons in the nucleus
of an atom of that element. Look at the diagrams of various atoms in
Fig. 32. What is the atomic number of each? the atomic weight?
How the electron theory explains isotopes. In 1815, William
Prout, a London physician, announced the theory that all the chem-
ical elements are made up of groups of hydrogen atoms only. Prout's
theory was not taken seriously for 100 years until Moseley's work
on atomic numbers made Prout's conclusion more plausible.
Since the nucleus of an atom of any element is composed of only
GROUP VIII
'Following lanthanum are 1 4 elements known as the rare earth elements (at. no. 58-71).
* * Following uranium are eight newly created elements of the actinide series (at. no. 93-100).
7
Nitrogen
14.008
15
Phosphorus
30.975
17
Chlorine
35.457
24
Chromium
52.01 34
Selenium
78.96
23
Vanadium
50.95 33
Arsenic
74.91
44
Ruthenium
101.7
45
Rhodium
102.91
46
Palladium
106.7
42
Molybdenum
95.95 52
Tellurium
127,61
51
Antimony
121.76
73
Tantalum
180.88 83
Bismuth
209.00
74
Wolfram
183.92 84
Polonium
210
75
Rhenium
186.31 85
Astatine
211
78
Platinum
195.23
92
Uranium
238.07
91
Protoactinium
231
TABLE 6. PERIODIC TABLE OF THE ELEMENTS
Henry 6. J. Moseley (1887-
1915), a pupil of Ernest
Rutherford, discovered the
law of atomic numbers in
1912. His brilliant career was
ended by his death at 27
during World War I.
neutrons and protons, and since the weight of each of these units
is really the weight of the hydrogen atom, it may have occurred to
you that the atomic weights of all the elements ought to be whole
numbers. But the fact that many atomic weights, for example, chlo-
rine (35.457) , are not whole numbers could riot be brought into
harmony with this idea.
In 1913, Theodore W. Richards found two different kinds of lead
with atomic weights of 206 and 207, respectively, and, in the same
year, two kinds of neon with different atomic weights were reported
also. The name isotopes was given to atoms of the same element hav-
ing the same chemical properties but different atomic weights. Dis-
coveries of isotopes of many other elements soon followed, one of
which (tin) is now known to have as many as 10 stable isotopes and,
hence, 10 different atomic weights.
The discovery of isotopes removed the obstacles to the acceptance
of Prout's idea. For elements, as we know them, are really mixtures
of isotopes having different atomic weights, each of which is a whole
number. Thus ordinary chlorine gas is really made up of some atoms
with an atomic weight of 35, other atoms of atomic weight 37, and
still others of atomic weight 39. Its accepted atomic weight, 35.457,
is the average of the atomic weights of the three different weights of
chlorine atoms in any sample of the gas. Here was another startling
discovery which helped to destroy Dalton's idea of an atom whose
atomic weight never changed.
The electron theory explains isotopes as caused by a different num-
ber of neutrons in each kind of atom. Thus, isotopes of chlorine with
atomic weights of 35, 37, and 39 behave alike chemically because
they have the same arrangement of planetary electrons. They differ
Fig. 31. Diagrams of the isotopes of chlorine.
Chlorine
isotope
35 '
164
Chlorine
isotope
37
Chlorine
isotope
39
ELECTRONS AND OTHER PARTICLES
165
in weight because of a difference in the number of neutrons in their
nuclei. We must, then, redefine the term element. A substance of
which all atoms have the same atomic number is an element.
How the electron theory explains valence. From the diagram of the
chlorine atom (Fig. 31) , you can see that its outermost ring contains
7 electrons. One more electron is needed to make the 8 electrons
needed to complete this ring.
An atom whose outermost ring is nearly complete has a tendency
to borrow enough electrons to complete this ring. An atom whose
outermost ring has few electrons tends to lose electrons. The number
of electrons gained or lost by an atom of an element is the valence of
that element. Since the chlorine atom needs to borrow only 1 elec-
tron to complete its outer ring, its valence is 1. In borrowing this
electron it becomes negatively charged and, as a result, the valence
of chlorine is negative. Hence, the valence of chlorine is —1.
An atom that lends electrons becomes positively charged. Hence,
elements whose atoms lend electrons have positive valences. Thus,
the sodium atom with an atomic number of 11 (roughly half of
22.997) may be pictured as shown in Fig. 32. As you see, the outer-
most ring contains 1 electron which the atom may lend. Hence the
valence of sodium is -f-1.
An atom whose outer ring is complete will neither lend nor bor-
row electrons. Elements whose atoms are of this type have a valence
of 0. The atom of neon is shown in Fig. 32.
Flectrons in the outermost ring of an atom, which may be either
borrowed or lent, are called valence electrons.
How the electron theory explains metals and nonmetals. A metal
is a lender of electrons. That is, the outermost ring of an atom of a
metal has less than four, or half the number (eight) required to
complete it. When such an atom lends electrons, it necessarily be-
comes positively charged. The valence of metals, therefore, is con-
sidered positive.
The atom of a nonmetal is a borrower of electrons. That is, its
outermost ring has more than four, or half the number of electrons
(eight) required to complete it. By borrowing electrons, such an
atom becomes negatively charged. The valence of nonmetals is,
therefore, negative.
Fig. 32. Diagrams of various atoms.
Sodium
Neon
Carbon
potassium
166
NEW WORLD OF CHEMISTRY
If the outermost ring of an atom of an element has just half the
number of electrons required to complete it, it may either borrow or
lend electrons. Such an element is said to be amphoteric. A common
example of an amphoteric element is carbon (atomic weight 12) ,
whose diagram appears in Fig. 32.
How the electron theory explains electric currents. According to
'modern theory, an electric current is a flow of electrons. Atoms of
metals, such as copper, silver, and gold, are good conductors of elec-
tricity, because some of their electrons are held loosely, and can move
freely through the solid. In general, nonmetals are poor conductors,
because their electrons are not held as loosely as those of metals.
How the electron theory explains chemical activity. An atom
tends to complete its outer ring of electrons. If an element such as
neon or argon already has its outer ring complete, that element is
inert. That is, its atom will not lend or borrow electrons, and hence
the element is completely inactive chemically. In general, the smaller
the number of electrons an atom must either borrow or lend to com-
plete its outer ring of electrons., the greater is the chemical activity
of that atom.
An atom, then, with either 1 or 7 electrons in its outermost ring
is extremely active. Such an atom is fluorine, whose atomic weight
is 19, and atomic number is 9. It has 7 electrons in its second ring
and will borrow 1 more electron. Potassium, whose atomic weight
is 39, has only 1 electron in its fourth ring, and hence can lend only
1 electron. Atoms such as those ot oxygen and magnesium have 2
electrons to borrow or lend and are quite active. Atoms of nitrogen
and aluminum have 3 electrons to borrow or lend and are not very
active. Generally, the farther away the outer ring of an atom is from
the nucleus, the less is the attraction of the nucleus for its electrons.
This helps explain the chemical behavior ot metals. For exam-
ple, potassium is more active than sodium since the electron in its
fourth ring can be lost more easily than the electron in the third
ring of sodium. Conversely, the closer the outer ring of an atom is
to the nucleus, the stronger is the attraction of the nucleus for its
electrons. This fact helps explain the chemical behavior of non-
metals. Fluorine is more active than iodine since its second ring has
Fig. 33. Chemical union of sodium and chlo-
rine according to the electron theory.
Sodium
Chlorine
sodium chloride
ELECTRONS AND OTHER PARTICLES 167
a greater attraction for the electrons of other atoms than does the fifth
ring of iodine. Both of these rules should be considered rough guides.
Many exceptions occur since the whole problem of chemical activity
is quite complex.
PRACTICE WORK ON THE ELECTRON STRUCTURE OF
ATOMS
1 . The at. wt. of sulfur is 32. Make a picture of its atom ac-
cording to the electron theory, and explain its valence and
chemical activity.
2. With the aid of a diagram, show why helium, at. wt. 4,
is inert.
3. With the aid ot diagrams, show why lithium, sodium,
and potassium belong to the same family ot elements.
4. What are the valence and chemical activity of an element
whose outer ring contains 4 electrons?
5. With a diagram, explain the valence of the (OH) radical.
How the electron theory explains chemical union and electro-
valence. Chemical activity is the tendency ot atoms to complete their
outer rings and form stable compounds. Chemical union is, there-
fore, the shifting or sharing of electrons in the outer electron rings
until a stable condition is reached.
A metallic atom with a valence of + 1 exhibits a strong attraction
tor a non metal lie atom whose outer ring needs 1 electron to com-
plete it. For example, the union of sodium, a very active metal with
a valence ot +1, with chlorine, a very active nonmetal with a valence
of — 1, may be represented thus:
The extra electron on the outer ring of the sodium atom shifts
over to the vacant space in the almost complete outer ring of the
chlorine atom. Now the two outer rings are both complete and the
resulting compound, sodium chloride, is very stable. Its water solu-
tion conducts electricity. Compounds formed by the shifting of single
electrons are polar or ionic compounds. Their valence is called elec-
trovalence.
Fig. 34. Formation of a molecule of fluorine
according to the electron theory.
shared pair of electrons
jf •
168 NEW WORLD OF CHEMISTRY
How the electron theory explains covalent compounds. In many
chemical reactions, there is no actual shifting of single electrons but
rather an equal sharing of a pair or pairs of electrons. The com-
pounds formed are nonpolar compounds and the valence is called
covalence. This kind of combination is generally very strong and the
molecules so formed hold together well. Nonpolar compounds are
stable and generally do not conduct electricity. Many organic com-
pounds, such as alcohol, glycerin, and sugar, are nonpolar.
The fluorine molecule (see Fig. 34) illustrates the sharing of a
pair of electrons. There is neither a gain nor a loss of electrons —
simply a sharing. Molecules of other gases that consist of 2 atoms,
such as oxygen and chlorine, exhibit this same sharing of electrons.
Polar and nonpolar compounds are discussed more fully on page 238.
How the electron theory explains oxidation and reduction. In
the equation CuO + H2 -^ Cu + H2O, copper oxide is reduced to
copper and hydrogen is oxidized to water. The valence of Cu has
changed from plus two (in CuO) to zero in free copper, and the cop-
per has gained 2 electrons (Cu++ + 2e -» Cu°) . The hydrogen has
lost an electron and changed to H+ in H2+O — (H° — e -» H+) .
Oxidation has been defined thus far as the union of a substance
with oxygen. Reduction, similarly, has been defined as the removal
of oxygen from a compound. From the viewpoint pf the electron
theory these terms take on much broader meanings. A loss of elec-
trons resulting in an increase in the valence of an element is called
oxidation. A gain of electrons resulting in a decrease in the valence
of an element is called reduction.
Oxidation and reduction in this broader sense need not involve
either oxygen or hydrogen. Thus, in the replacement of the iodine
in potassium iodide by chlorine, as explained in the preceding chap-
ter, we have an oxidation-reduction reaction. According to the elec-
tron theory, this reaction is:
K+I- + Cl° -> K+C1- + 1°
The iodine in KI has lost an electron and changed to free iodine, 1°.
It has been oxidized by chlorine which has gained an electron and
changed from free chlorine to the negative chloride ion., Cl~ (see
page 232) .
How the electron theory aids in balancing equations. We may see
from this oxidation-reduction equation how each atom of iodine lost
an electron and was oxidized and how each free chlorine atom gained
an electron, forming a chloride ion with a negative charge. This was
a relatively simple reaction. But how is a more complex oxidation-
ELECTRONS AND OTHER PARTICLES 169
reduction equation balanced and how does our understanding of the
electron theory help us to find correct coefficients to use in balanc-
ing the equation? Let us consider some examples:
EXAMPLE A: It is desired to reduce ferric chloride, FeCl3, to
ferrous chloride, Fed,, by the use of the evil-smelling gas,
H2S — a common reducing agent. In the reaction, free sulfur
and hydrochloric acid are formed also.
1) Write the unbalanced equation for the reaction.
FeCl8 + H2S -» FeCl2 + S + HC1
2) Select the atoms which, according to the electron theory,
are either reduced or oxidized, that is, either gain or lose elec-
trons. Place the valences involved at the upper right of the
symbol of each atom.
Fe+++Cl3 + H2S— -> Fe++Cl2 + S° + HC1
3) By inspection, we see that each atom of iron gains one
electron and is reduced; each atom of sulfur loses two electrons
and is oxidized. The number of electrons involved in the oxi-
dation of H2S is twice as great as the number of electrons in-
volved in the reduction of one molecule of FeCl3. Therefore,
twice as many molecules of FeCl3 must have been reduced.
Indicate this by writing the coefficient 2 before FeCl3 in the
equation. Then balance the equation by the methods you have
already learned.
2FeCl8 + H2S -» 2FeCl2 + S + 2HC1
In actual practice, chemists usually consider only the changes in
valence of the atoms involved rather than the shifting or sharing of
electrons. The sum of the changes in valence on both sides of the
arrow must be equal. Consequently, this method of balancing equa-
tions is known as the valence-change method.
EXAMPLE B: Using the valence-change method, write the equa-
tion for the reaction between the oxidizing agent, potassium
permanganate, and hydrochloric acid. The products are potas-
sium chloride, manganese chloride, water, and free chlorine.
1) Write the unbalanced equation, omitting the subscript
from the free chlorine. It is not necessary to indicate the
valence of those atoms which retain the same valence through-
out the reaction.
4 + HC1- -» KC1 4- Mn++Cl2 + H2O + Cl°
170 NEW WORLD OF CHEMISTRY
2) Note that the following changes occur:
a) Each atom of manganese in KMnO4 gains five elec-
trons in combining with chlorine to form MnCl2; therefore the
valence loss of the manganese is five.
(b) Each atom of chlorine loses one electron in becoming
free chlorine; therefore the valence gain of each atom of lib-
erated chlorine is one.
3) The five valences lost by the manganese must be balanced
by the valence gained by the chlorine since no other element
changes valence in the reaction. To do this, we must show that
for each molecule ot KMnO4 used, five atoms of Cl are lib-
erated. Therefore, write the coefficient 5 before the free Cl:
KMn04 + HC1 -> KC1 + MnCl2 + H2O + 5C1
4) As we know, free chlorine exists as a molecule composed
of two atoms, therefore we add the subscript, 2, to Cl. In order
to keep the valence changes in balance, we consider that two
molecules of KMnO4 are used in liberating five molecules of
C12. Therefore, write the coefficient 2 before KMnO4.
2KMn04 + HC1 -> KC1 + MnCl* + H2O + 5C12 f
5) Complete the balancing of the equation in the usual
manner. Check your work carefully.
2KMnO4 + 16HC1 -> 2KC1 + 2MnCl2 + 8H2O + 5C12 T
From this discussion, we can formulate the following method of
balancing oxidation-reduction equations:
1) Write the unbalanced equation.
2) Write the changes in valence at the upper right of the
symbols of the atoms that are oxidized and those that are re-
duced.
3) Find the valence gains and losses and decide on the coef-
ficients that will make them equal. Remember that molecules
composed of two atoms may be involved and adjust the coeffi-
cients accordingly.
4) Complete the balancing of the equation by the usual
method.
Changing theories and the spirit of modern science. In 1924, the
electron theory, backed by a mass of experimental evidence, was
quite generally accepted. In that year, Prince de Broglie (de bro'y')
suggested that the electron is not merely a particle of electricity, as
the electron theory explained, but like light, is composed of, possesses,
or perhaps is attended by, a group of waves.
ELECTRONS AND OTHER PARTICLES 171
By 1927, two Americans, Davisson and Germer, proved de
Broglie's theory experimentally, showing that both electrons and
protons possess a property of light (a wave phenomenon) . In 1929,
de Broglie was honored with the Nobel prize in physics for his theory.
The theory that matter possesses wave properties created an up-
heaval in the existing theories. However, we still find it convenient
to regard the electron, proton, and neutron as tiny individual par-
ticles. If new facts show we are wrong, we shall scrap this concept.
In every phase of science, this practice is followed. Old ideas are
retained as long as they are useful. They may be altered somewhat
to fit newly-discovered facts. However, if enough new data are accu-
mulated to prove them incorrect, the old theories are abandoned.
It is the belief of scientists that tradition should never be allowed to
stand in the way of greater enlightenment.
Albert Einstein, one of the most eminent of living scientists, when
he was suddenly confronted with new facts which could not fit his
own theories, expressed the spirit of science thus: "The new facts
have smashed my old ideas like a hammer blowl" And he went on
to change some of his most cherished theories.
YOU WILL ENJOY READING
Jaffe, Bernard. Crucibles: The Story of Chemistry, pp. 289-
311. Simon and Schuster, New York, 1948. A sketch of Henry
Moseley, whose lifework was done in 4 short years. Before the
world knew this genius, he died.
Moulton, Forest R., and Schifferes, Justus J., Ed. The Auto-
biography of Science, pp. 502-509. Doubleday Doran & Co.,
New York, 1945. Contains the original papers dealing with the
origin of the electron theory.
USEFUL IDEAS DEVELOPED
1 . Electric charges caused by rubbing or friction are called
static electricity.
2. Electrons are tiny particles of negative electricity. A
cathode ray is a stream of electrons. Our knowledge of electrons
is the result of the work of William Crookes, J. J. Thomson,
Robert A. Millikan, and others.
3. Protons are particles oi positive electricity. A proton is
1837 times heavier than an electron. It was discovered and
named by Ernest Rutherford.
172 NEW WORLD OF CHEMISTRY
4. A neutron is an electrically neutral particle composed of
one electron and one proton. Its mass is about the same as
that of a proton, or of the atom of the lightest hydrogen
isotope.
5. All atoms are composed of electrons, protons, and neu-
trons with the exception of the atom of the lightest hydrogen
isotope which is composed of one electron and one proton.
Protons and neutrons are located in the nucleus of the atom.
Electrons revolve about the nucleus in orbits. The relative
distance between the nucleus and the planetary electrons is
so great that we may say the atoms of all elements are largely
empty space.
6. A fixed number of electrons is required to complete
each electronic orbit of the atom. The first ring is complete
when it contains 2 electrons; the second, when it contains 8;
the third, when it contains 18; the fourth, when it contains 32.
The outermost ring never contains more than 8 electrons.
7. The atomic number of any element is the number of
free protons in the nucleus of its atom. According to the law
of atomic numbers, the elements can be arranged in a periodic
table in the order of their increasing atomic numbers.
8. The atomic weight of any element is the sum of the
protons and neutrons in the nucleus of one of its atoms.
9. Isotopes are different forms of the s^me element. They
possess the same chemical properties, but have slightly different
physical properties. Isotopes of an element all have the same
atomic number, but different atomic weights.
10. The valence of an element is the number of electrons
that its atom must borrow or lend to complete its outermost
ring. Electrons in the outermost ring of an atom which may
be borrowed or lent are valence electrons.
11. An element is metallic if the outer ring of its atom con-
tains less than half the number of electrons necessary to com-
plete the ring. In other words, metals are lenders of electrons.
12. An element is nonmetallic if the outer ring of its atom
contains more than half the number of electrons necessary to
complete the ring. In other words, nonmetals are borrowers
of electrons.
13. An element is amphoteric if the outer ring of its atom
contains exactly half the number of electrons necessary to com-
plete this outer ring.
14. An electric current is a flow of electrons.
15. Chemical activity depends upon the number of elec-
trons in the outermost ring of the atom of an element. When
the outermost ring is complete, the element is inert; if the
outermost ring lends or borrows one electron, the element is
ELECTRONS AND OTHER PARTICLES 173
very active; if the outermost ring lends or borrows three elec-
trons, the element is not very active.
16. Chemical union is the shifting or sharing of the electrons
in the outer rings until a stable condition is reached.
17. Electrovalent compounds are those formed by the shift-
ing of single electrons; covalent compounds are those formed
by the sharing of a pair or pairs of electrons.
18. Oxidation, a loss of electrons, increases the valence of
an element. Reduction, a gain of electrons, decreases the
valence of an element.
19. New theories are constantly being advanced about the
nature of matter. Recently it has been suggested that both the
proton and electron possess wave properties. It is necessary for
scientists to be ready to give up theories when facts show that
newer theories are more nearly accurate.
USING WHAT YOU HAVE LEARNED
Group A
1. (a) How may static electricity be produced? (b) How
did the discovery of electrons explain this phenomenon?
2. (a) Who discovered the electron? (b) How did his dis-
covery explain the behavior of cathode rays? (c) Who deter-
mined the mass of the electron? (d) What were his findings?
3. (a) Who discovered the proton? (b) What principle
of electricity led him to his discovery? (c) Do the protons of
different elements differ?
4. (a) What is the neutron? (b) What scientists took part
in its discovery? (c) Why is the neutron considered approxi-
mately equal in mass to the proton?
5. (a) Describe the general arrangement of protons, elec-
trons, and neutrons in an electrically neutral atom, (b) De-
scribe the atom in terms of the solar system.
t . . .
6. Explain how elements differ according to the electron
theory.
7. (a) Who developed the first periodic table of atomic
weights? (b) What was learned from arranging the elements
in such a table?
8. (a) What relationship between atomic number and pro-
tons was demonstrated by Moseley? (b) Why is Moseley's
periodic table more fundamental than Mendeleyeff s?
174 NEW WORLD OF CHEMISTRY
9. (a) For the first 17 elements, what is the general re-
lationship between atomic number and atomic weight? (b) Il-
lustrate your answer.
10. (a) Define atomic weight in terms of the electron the-
ory, (b) Why do most elements have atomic weights which are
not whole numbers?
11. (a) What is an isotope? (b) Explain why all isotopes
of an element have the same chemical properties.
12. (a) How does the electron theory explain valence?
(b) How does it explain the behavior of metals? (c) of non-
metals? (d) of amphoteric elements?
13. Explain the relationship between the number of elec-
trons in the outer ring of an atom and the chemical activity
of that element.
14. (a) Make a diagram of the structure of the element
whose atomic number is 20. (b) Describe some of the element's
chemical properties.
15. Phosphorus (P) has an atomic weight of 30.98. (a) Dia-
gram the structure of its atom, (b) Describe its chemical
nature.
16. An atom has a nucleus containing 1? protons and 18
neutrons, (a) Make a structural diagram of the atom, (b) De-
scribe its chemical properties.
17. The atomic weight ot beryllium (Be) is 9.02. (a) What
are the atomic number, valence, and chemical properties of
beryllium? (b) Is it a metal or nonmetal?
18. The atomic weight of curium, one of the newly-discov-
ered elements, is 242; its atomic number is 96. How many
neutrons are in the nucleus of one of its atoms?
19. Make a diagram of (a) the oxygen atom and (b) the
sulfur atom, (c) Explain, in terms of the electron theory, why
they resemble one another chemically.
20. (a) Draw a diagram of the structure of a molecule of
potassium bromide, KBr. (b) Using your knowledge of the
electron theory, explain why the two atoms unite.
21. What is the difference between electrovalence and co-
valence?
ELECTRONS AND OTHER PARTICLES 175
22. How does the electron theory explain why two atoms of
hydrogen unite with one atom of oxygen to form one molecule
of water?
23. How does the electron theory explain the fact that cer-
tain elements are more active chemically than others?
24. In terms of the electron theory, explain why some com-
pounds are stable and others unstable.
25. Explain oxidation and reduction in terms of the electron
theory. Use the equation for the reducing action of H2 on CuO
to make your answer clearer.
Group B
26. Study the theories Benjamin Franklin held regarding
the nature of electricity. Compare his views with the modern
theory.
27. (a) What is a cloud chamber? (b) For what purpose is
it used? (c) What are alpha particles? (d) fog tracks?
28. What evidence is there that protons are always found
inside the nucleus of an atom?
29. Using the valence-change method, balance the follow-
ing equations:
a) HC1 + Mn02 -» MnCl2 + H2O + C12 T
b) C12 + H20 -> 2HC1 + O T
c) KC103 -> KC1 + 02 T
d) S + HNO3 -» H2SO4 + NO |
30. Why does carbon have the 2 valences, -f-2 and -f-4?
31. Explain why the valence of a free element must be zero.
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Harkins predicted the discovery of the neutron, and
Mendeleyeff predicted the discovery of several elements. Both
predictions were verified. Such incidents in the history of sci-
ence are not as rare as in other fields such as economics and
politics. Give reasons for this.
2. Construct a model of the sodium atom based on the elec-
tron theory. You may use wire for the orbits, a small ball for
the nucleus, and copper rods for the electrons.
NUCLEAR ENERGY
AT LAST!
The United States knows that peace-
ful power from atomic energy is
no dream of the future. That capa-
bility, already proved, is here now —
today. I propose an atomic energy
agency under the United Nations to
apply atomic energy to the needs of
agriculture, medicine., and other
peaceful activities . . . to provide
abundant electrical energy in the
power-starved areas of the world.
Dwight D. Eisenhower, December 8,
The dream of harnessing the energy of the atom. The atomic
bomb dropped on Japan on August 6, 1945, was one of those incred-
ible pinnacles toward which fate drives unsuspecting man. Thou-
sands of scientists who had studied the atom and discovered the
electron, proton, and neutron had little idea of the genii they were
uncorking. Chiefly, they were men and women trying to learn more
about the nature of matter, each concerned with his own particular
problem.
A few scientists, it is true, saw the possibility of someday learning
enough about the atom to enable man to release the tremendous
forces locked within it. If it could be done, the energy locked within
a single lump of coal would be enough to drive a huge ocean liner
around the world. As scientists, they hoped for an achievement which
they believed to be even greater than the discovery of electricity.
They dreamed of opening a door to an age of limitless power and
thus lifting the standards of living of all the peoples of the world.
When the news of triumph finally came, it surprised even the most
optimistic of scientists. The great marvel, said President Truman,
". . . is not the size of the enterprise, its secrecy or its cost, but the
176
NUCLEAR ENERGY AT LAST!
177
achievement of scientific brains in putting together infinitely com-
plex pieces of knowledge held by many men in different fields of sci-
ence into a workable plan." The story of some of these men and the
bits of knowledge they accumulated has been told in earlier chap-
ters. More of the story can now be told.
An amazing scientific discovery startles the world. Late in 1895,
William Roentgen (rimt'ge?n) was working in a dark room with a
Crookes tube covered with black paper. He noticed that while an
electric discharge was passing through the tube, a small screen cov-
ered with a chemical, barium platinocyanide, lying on a table several
feet away, gave off a strange glow.
This was curious and unusual behavior. He thought it must be
caused by rays, powerful enough to penetrate not only the glass of
the Crookes tube but several feet of air as well. He tested the pene-
. trating power of these strange rays by producing them in front of
several objects of varying hardness, including a hand behind which
he had placed a sensitized photographic plate. To his astonishment,
the film when printed showed a hand with the bones much darker
than the surrounding flesh. He had, so to speak, taken pictures
through an opaque solid, a truly remarkable feat!
These rays in some ways acted like light, but differed from light
in being of much shorter wave lengths which could pass through
even solid objects. He named these rays x-rays. X-rays are produced
by the bombardment of matter by a stream of rapidly moving elec-
trons, or cathode rays. In the cathode tube, as the cathode rays strike
the anode, they release a small part of the energy as x-rays.
X-ray machines (right) are used in industry to detect
flaws inside metal castings like the crankshaft (left).
ml Klrrtr,
178 NEW WORLD OF CHEMISTRY
Another accident, and the curtain rises on the drama of radio-
activity. Soon after the discovery of x-rays, another accident occurred
in the laboratory of Henri Becquerel (bek-reT) . He was testing the
effect of sunlight on various ores, among which he had, fortunately,
included an ore of uranium.
Quite by chance he placed a piece of the ore containing the ele-
ment uranium upon a fresh photographic plate enclosed in a light-
tight envelope lying on a table in his darkroom. When he examined
this plate, he found that it had been changed under the very spot
on which the ore had rested. This was not the sort of accident to
reach the front page of the newspapers, as the discovery of x-rays had
done. But its results were tremendously important.
A new world of radioactivity is discovered. Becquerel could not
explain what had happened. He repeated the experiment with other
ores containing uranium. Pitchblende, he found, emitted similar
rays, which affected a photographic plate even more than the other
uranium compounds. He suspected some unknown element to be
the cause, and asked Madame Marie Curie (ku-re') , a Polish girl
working as a science teacher and research worker in Paris, to under-
take the isolation of this unknown element.
Madame Curie and her husband, Pierre, set out to track down the
cause of this peculiar behavior of pitchblende. They boiled and
cooked a ton of this ore, sent to them from the pitchblende mines of
Austria; they filtered and separated out impurity after impurity.
Years of almost endless work passed, and, though they labored under
extreme difficulties, Madame Curie wrote years later: "It was in that
Marie Curie in her laboratory. For her
isolation of radium and polonium,
Madame Curie was awarded the
Nobel prize for chemistry in 191 1.
NUCLEAR ENERGY AT LAST!
179
miserable shed of a laboratory that we passed the best and happiest
years of our lives." Finally in 1898 they succeeded in obtaining a few
crystals of a salt of a new element, which they named radium.
The new discovery was made public. A strange element had been
discovered by a woman. Its salts shine in the dark like tiny electric-
light bulbs and emit heat continuously. This element is a powerful
poison — even acting from a distance. It causes severe burns, and has
brought premature death to a number of scientists who have han-
dled it for long periods. It makes the air around it a good conductor
of electricity.
Because radium captured the imagination of the entire world, its
discovery was a great stimulus to further research. This led to the dis-
covery that a number of other elements resemble radium in their
ability to break down and emit several kinds of rays. This property,
called radioactivity,, is possessed by thorium, uranium, polonium,
radon, and several other elements.
In 1902, Rutherford and Soddy, another English scientist, ex-
plained the disintegration or breaking down of radium. Atoms of
radioactive elements, they said, are not stable. They explode spon-
taneously, giving off three types of rays: alpha and beta particles and
gamma rays. The gamma rays are similar to x-rays; the beta rays are
electrons. Rutherford later showed that the alpha rays are electrified
particles consisting of nuclei of helium atoms. The disintegration or
decay of radium is indicated in Table 7.
What is the half -life of an element? The half -life of an element is
the time required for the radioactivity of a given amount of an ele-
ment to decay to half its original value. For example, starting with
one gram of radium, it takes 1620 years for half of it to change to
lead. At the end of the next 1620 years, half of the 0.5 gram which
is left changes to lead, and this process continues at the same rate.
TABLE 7.
DISINTEGRATION
OF RADIUM
1 RADIUM A
at. wt. 21 8
half-life 3 min.
loses one helium
nucleus of
at. wt, 4 and
change
at. wt. 214
half -life f9.5 min.'
loses one helium
nucleus of
at. wt. 4 and
changes
| POLONIUM
at. wt. 210
half-life 140 days
loses one helium
nucleus of
at. wt. 4 and
changes
at. wt. 206
a stable element
which is the end
product of radium
disintegration
~-!^ radiation
,90$ molecule
metal cylinder (-)
tungsten wire ( + ) 9|ass fube
to high potential
to R and detector *
Fig. 35. Geiger counter. An entering electron pro-
duces a discharge, causing current to flow in the
high resistance R until the fall of potential across
R reduces the potential and discharge stops. The tube
is then ready for the arrival of a second electron,
(left) Use of the Geiger counter to test uranium ore
for the amount of radiation.
Standard Oil Company (XJ.)
The Geiger counter, detector of nuclear disintegration products.
One type of Geiger counter consists of an argon-filled tube contain-
ing a metal cylinder and a thin wire. Between these two electrodes,
a very high voltage is maintained and so adjusted that the tube is on
the verge of discharge. When an electron enters the tube and col-
lides with a gas molecule, the tube is discharged and a How of cur-
rent is produced. This current operates a headphone, or produces a
loud click or flash of light. The greater the radiation, or the closer
the tube is to the nuclear disintegration, the greater the effect.
Alpha particles are also detected by the Geiger counter. Neutrons,
which have no charge, are detected indirectly with the use of a tube
which contains a boron compound. The boron nucleus absorbs neu-
trons and produces particles which may be detected. Gamma rays
are detected by the secondary electrons they produce.
Alpha particles, which travel at 10,000-20,000 mi. /sec., have the
least penetrating ability. About five centimeters ot air, a sheet of
paper, or a thin sheet (0.1 mm.) of aluminum will stop them. Beta
particles (electrons) , liberated at speeds more than six times that of
alpha particles, require several meters of air and several millimeters
of aluminum to absorb them. Gamma rays have still greater pene-
Fig. 36. Alpha, beta, and gamma rays
have different penetrating powers and
are affected differently by a magnetic field.
radium . , , v
salt. alpha (a)
^J_ particles I
^Hl .....*.. JW"ii
•• beta (fiTT"^
lead Particles
box m°9net
aluminum
•( foil
aluminum
gamma
(7)
rays
NUCLEAR ENERGY AT LAST!
181
trating power — several centimeters of aluminum are required to
stop them. Fast neutrons have the greatest penetrating power. All
produce burns. Radiation absorption is measured in roentgens (r.) .
Uses of radium. Radium, and radon gas sealed in tubes, are used
in treating skin diseases and cancer.
Considerable quantities of radium are used in the detection of
(laws in castings, forgings, and welds. Parts for aircraft and turbine
casings are among the many kinds of equipment tested in this way.
The method of testing is simple. From 25 to 1000 milligrams of
radium sulfate are placed in the center of a circle of articles to be
tested. X-ray films are placed on the backs of the specimens. The
penetrating gamma rays from the radium salt produce a shadow-
graph on the film, quite like the kind obtained with x-rays. Defects
as small as 0.25 percent of the thickness of the article can be seen
clearly.
The price of radium is about $20,000 a gram. Carnotite, an ore of
both radium and uranium, is found in Utah and Colorado. The rich
deposits of pitchblende discovered on the shore of remote Great
Bear Lake in northern Canada in 1931 compete on a favorable basis
with the huge deposits in the Belgian Congo which constitute our
most important source of uranium.
The attack upon the nucleus of the atom. The disintegration of
radioactive elements indicates that the building blocks of the atom's
nucleus consist of neutrons and free protons, both of which can be
emitted. Helium nuclei, which are liberated also by radioactive ele-
ments, are themselves composed of neutrons and free protons.
In the effort to learn more about the nucleus, scientists began
bombarding atoms with various kinds of swiftly moving projectiles.
One of the most important machines for acceleration of particles is
the cyclotron developed in 1929 by Ernest O. Lawrence. Some other
Brookhavcn National Labti
The magnet in this giant cos-
matron at Brookhaven has
an inside diameter of over
60 feet and weighs 2200 tons.
182
NEW WORLD OF CHEMISTRY
accelerators for smashing atoms are die betatron, which * accelerates
electrons, and the synchroton, cosmotron, and bevatron, all of which
speed up protons to even greater speeds than the cyclotron.
During these researches, several discoveries were made — the posi-
tron, artificial radioactivity, and mesons. The positron is an ex-
tremely short-lived particle, having the same mass as the electron
but opposite charge. It was discovered and named by Carl D. An-
derson, of the California Institute of Technology, in 1932. When a
positron reacts with an electron, both particles disappear and gamma
radiation is formed.
In the nucleus of the atom a tremendous energy lies like a coiled
spring caused, in part, by the many protons which repel each other
because they are all of the same electric charge. Mesons may be the
binding energies that hold the protons in the nucleus together.
There are several types of mesons. The heavier meson (pi-meson) is
produced when a nucleus is broken up during a collision. The
lighter meson (raw-meson) is a decay product of the heavy meson.
Mesons are very short-lived — less than a millionth of a second.
Summary: Units of matter thus far identified. Let us pause tor a
few minutes to list these units of matter, their mass, charge, and
other related items of: information.
'«*>''.; ** * ti^M'^ <a
*-.•• ./;.|i
\4t
" ^&«*""
y.$*e
V»***» -
t NAME
'j^wr""*"* <•-* >*~
Hfc*:* * •
MASS
CHARGE
WHERE
FOUND
HOW OBTAINED
OR FORMED
'• ^ .*'•>• t 1
DISCOVERED BY
AND WHEN
» Electron
J (6)
1/1 837 that of the
hydrogen atom
Negative
Outside
nucleus
In Crookes tube, as
cathode rays
Joseph John
Thomson, 1 897
& Proton
4»+>
Approximately that
of hydrogen
Positive
Inside
nucleus
Stripping an electron
from hydrogen
Ernest Rutherford,
1911
2 Neutron
1 (°nl>
Approximately that
of hydrogen
Neutral
Inside
nucleus
Bombarding beryllium
with helium
James Chadwick,
1932
1 Positron
I Ut«°)
Same as the
electron
Positive
Inside
nucleus
Radioactive nitrogen
disintegration
Carl D. Anderson,
1932
1 Meson*
E8. |
? I
Heavy, or pi, meson
is about 285 times
that of electron.
Light, or mu, meson
is about 21 5 times
that of electron.
Positive,
negative,
and
neutral
Inside
nucleus; as
secondary
cosmic rays
in upper
atmosphere
Bombarding atoms
with primary cosmic
rays (protons) or with
helium nuclei of 300
million or more
electron-volts in
cyclotron
Carl D. Anderson
and
S. Neddermeyer
discovered the
mu-meson in 1 937
andC. F. Powell
the pi-meson in
1947.
5 'This particle was predicted by H. Yukawa, a Japanese, in 1934. It was named meson,
s meaning Intermediate porfic/e It changes into something else in less than a millionth
$ of ? second, and it is supposed to travel at a speed nearly that of light.
NUCLEAR ENERGY AT LAST! 183
Ancient and modern alchemy. As you know, modern chemistry
sprang from alchemy, which was practiced for more than 20 centu-
ries. The chief goals of the alchemists were to change the base metals,
such as lead and iron, into gold and to find an elixir that would cure
all disease. Although among the alchemists there were many honest
enthusiasts, the annals of their queer practice are filled with accounts
of liars and charlatans. In many museums of Europe we can still see
shiny, yellow metals reputed to be gold, made by the deceptions of
the * 'gold-cooks" of European courts.
Today the alchemists' dream of changing one element into an-
other, called transmutation, has come true. Radium changes, of its
own accord, into helium, lead, and other elements. Besides this natu-
ral transmutation, however, chemists have succeeded in artificially
transmuting many nonradioactive elements. The first such transmu-
tation was achieved in 1919 by Rutherford, who bombarded nitro-
gen with helium nuclei. The nucleus of the nitrogen atom was
changed, and one proton was liberated, the remaining nucleus be-
coming the heavy isotope of oxygen with atomic weight 17.
How nuclear reactions are written. Transmutation or nuclear re-
actions are like chemical equations and must balance. The sum of
the subscripts must be the same on both sides, and so must the sum
of the superscripts.
Rutherford's transmutation may be expressed as follows.
The subscripts (7 + 2-»8+l) represent atomic numbers and the
superscripts (14 + 4— » 17 + 1) atomic weights.
Recent experiments indicate that other elements may be built
up from lighter elements, for example, carbon from beryllium (a rare
metal lighter than aluminum) . This transmutation, (Fig. 37) , may
be used to produce high-speed neutrons by mixing powdered beryl-
lium with a trace of a radium salt which emits helium particles.
Transmutation, as you see, involves changes in the nuclei of atoms
rather than in the shifting or sharing of electrons which produces
only chemical changes.
Fig. 37. Transmutation of beryllium into carbon.
Beryllium + Helium ^ Carbon
(at. wt. (at. wt.
9)
184 NEW WORLD OF CHEMISTRY
The search for the key to nuclear energy. Albert Einstein, in 1905,
advanced the idea that matter and energy were really different forms
of the same thing, and that matter could be changed into energy —
at least theoretically. He developed a mathematical equation to ex-
press the conversion of matter into energy:
E = mc2
where E is energy expressed in ergs,* ra is mass expressed in grams,
and c is the speed of light, expressed in centimeters per second.
According to this matter-energy conversion equation, one pound
of matter (for example, one pound of coal or of uranium) is equiva-
lent to about 11 billion kilowatthours, if completely changed into
energy. This is about two and one-half times the amount of electric
energy produced in an entire year by the largest steam-electric gen-
erating plant in the country. In burning the same amount of coal,
approximately four kilowatthours of energy are obtained. In terms
of energy produced, oxidation is, therefore, an extremely inefficient
process.
These ideas, of course, were all theory. However, a bit of confirma-
tion came in 1932. In that year, Cockcroft and Walton, working in
Rutherford's laboratory, bombarded lithium with high-speed pro-
tons produced by accelerating hydrogen nuclei by Yneans of high
voltages. They obtained helium (alpha particles) with energies al-
most 100 times as great as the energy that was used to break the
lithium atom. This extra energy comes from the conversion of mat-
ter into energy, in accordance with the equation formulated by Ein-
stein, thus:
Lithium + hydrogen — > 2 helium + energy
3Li7 + iH1 — » 22He4 + 600,000 electron-volts
Mass 7.0180 + mass 1.0076 -> mass 2(4.0029) f
8.0256 -> 8.0058
Approximate loss of mass = 0.02
* 1 erg == 1/980 gm.-cm. of work = approximately the energy required to lift
a postage stamp to a height equal to its thickness. The speed of light is
186,000 mi./sec. or 30,000,000,000 cm./sec.
f Note that the atomic weight of the isotope of the lithium used differs from
the atomic weight given in the table on page 162, which is an average of the
atomic weights of all the isotopes of lithium. The hydrogen here refers to the
proton, which is slightly lighter than the hydrogen atom, whose weight is given
in the table on page 162. The weights given in the table on page 162 have been
rounded off to three decimal places; hence, helium is shown there as 4.003 instead
of 4.0029 as in this equation.
NUCLEAR ENERGY AT LAST! 185
However, the method used by these experimenters was not efficient,
and there was no great excitement over their news.
The key is found. In the meantime, other scientists were working
in this same field. In 1934, a young Italian physicist, Enrico Fermi
(far'mi) , who later left fascist Italy to become professor of physics
at Columbia University, bombarded uranium with neutrons and
thought he had created a new element No. 93. Then four years later
Otto Hahn and F. Strassman repeated Fermi's experiment in Berlin.
They bombarded uranium with slow neutrons and, instead of pro-
ducing a new artificial element, they obtained two other natural ele-
ments and a great deal of energy. Unable to explain what had hap-
pened, Hahn and Strassman nevertheless published their findings.
Lise Meitner (mlt'ner) , an eminent woman scientist working with
Hahn, interpreted the results and passed the information on to Niels
Bohr, Nobel prize winner in physics, in Copenhagen. Dr. Meitner
was forced to flee Germany by the Na/is. Dr. Meitner believed that
when uranium is bombarded by slow neutrons, the atom of uranium
actually splits by a process called nuclear fission, forming barium and
krypton. But what is even more important, great quantities of en-
ergy are released, perhaps as much as 11 million kilowatthours per
pound of uranium. And, this is only a small part of the energy that
would be produced if all the uranium were converted into energy.
The stage is set. Very soon after, a most important conference
was held in Washington, D.C. Atomic physicists from American col-
leges and famous scientists from foreign nations were present. Niels
Bohr was there, and so was Enrico Fermi. At this meeting, Bohr and
Fermi discussed the ideas of Meitner. Bohr suggested that it was the
U-235 in the uranium that actually split. Fermi suggested that, in
the fission of uranium by neutrons, other neutrons might be emitted.
These emitted neutrons could attack other uranium atoms. If this
were true, the possibility of a chain, or self-propagating, reaction
that would unlock the door to nuclear energy was near at last.
Brown Ilroflu
Enrico Fermi (1901-1954), winner of the 1938
Nobel prize for physics, played an important
role in our government's nuclear research
program both during and after World War II.
186 NEW WORLD OF CHEMISTRY
Before the meeting in Washington was over, experiments to con-
firm nuclear fission had begun, and confirmation of the emission of
neutrons was soon obtained. By midsummer of 1940, the important
facts regarding nuclear fission had been discovered and were known
by many scientists. And although a chain reaction had not been ob-
tained, its possibility was clear and several methods of producing
it had been suggested. Then suddenly, World War II clamped tight
the door of censorship on all research relating to the release of
nuclear energy. For five years, the outside world was kept in the dark.
Nuclear energy unleashed! With the sudden dropping of the first
atomic bomb on Hiroshima in August, 1945, the veil was partly
lifted on research on nuclear fission and the production of chain
reactions. Early in 1940, Franklin D. Roosevelt and Winston
Churchill had pooled the efforts of British and American scientists
on a research program, the like of which the world had never seen.
Its goal was the release of nuclear energy for the production of a
weapon with which to win the war against the Axis nations more
quickly. Knowledge that research on such a weapon was being car-
ried on in Nazi laboratories compelled quick, cooperative action.
The race was on — the prize, the world. The United States gov-
ernment invested two billion dollars in ". . . the greatest scientific
gamble in history — and won."
The term atomic energy has long been used to describe the tre-
mendous power which is released when nuclear fission occurs. How-
ever, the term nuclear energy is now preferred since it is more truly
descriptive of the processes involved.
A chain reaction from U— 235. The first controlled chain reaction
was achieved on December 2, 1942, at the University of Chicago. The
fissionable material used was pure U-235 obtained from natural
uranium ores containing a mixture of three isotopes: U-234, U-235,
and U-238. Even though only about one part in 140 of this mixture
is pure U-235, this isotope is used because it is most susceptible to
nuclear, fission by slow neutrons. What happens in the nuclear fission
of U-235 may be represented as:
U-235 + neutron — - » Ba + Kr + 2 or 3 neutrons + energy
At Oak Ridge, Tennessee, U-235 was laboriously separated from its
other isotopes by an electromagnetic method. A compound of ura-
nium, UF6, was passed in the form of a gas between the poles of a
magnet. The lighter isotope, U-235, was deflected more than its
heavier partners and thus separated.
penetrating
radiation
liberated neutrons energy
200,000,000
electron-volts
56 +
82 n
Fig. 38. A possible chain reaction.
A newly created element, plutonium, for the A-bomb. Few details
about the A-bombs exploded over Hiroshima and Nagasaki have
been released. It is definitely known, however, that two fissionable
elements were produced for use in bombs; namely, U-235 and plu-
tonium. Plutonium, which has fission properties similar to U-235, is
a newly created element of atomic number 94. It was named for the
planet Pluto, which lies beyond Uranus in the solar system. Together
with another new element, neptunium of atomic number 93, plu-
tonium was first prepared in 1940 with the aid of Lawrence's cyclo-
tron by E. M. McMillan and P. H. Abelson.* These were momentous
discoveries.
The formation of neptunium and plutonium may be represented
by the following three equations:
1) Uranium 238+ neutron
uranium 239
23 mirv
2) Uranium 239 ^-neptunium 239 + electron
* Traces of these two elements were later found in some uranium ores.
In 1945, elements 95, americium (Am) , and 96, curium (Cm) , were obtained
by bombarding plutonium and uranium with swiftly moving helium nuclei.
Then came elements 97, berkelium (Bk) , and 98, californium (Cf) , in 1949
and 1950. These elements were named after Berkeley and California, the city and
state in which they were first produced. Elements 99 and 100 were created in 1954.
187
188 NEW WORLD OF CHEMISTRY
This change occurs by the breaking down of one neutron in the nu-
cleus of U-239 into one proton and one electron, which escapes.
2.3 days
3) Neptunium 239 *> plutonium 239 + electron
This change occurs by the breaking down of one neutron in the
neptunium 239 nucleus into one proton and one electron, which
escapes.
Plutonium, in turn, becomes U-235 by natural radioactive disinte-
gration, or more rapidly by the action of slow neutrons to which it
is extremely sensitive. The change is indicated by the following nu-
clear equation:
Plutonium — » U-235 -f helium
What is meant by "critical size"? An A-bomb is set off by suddenly
bringing together two separate blocks of fissionable material, each
of which is smaller than the critical size, but which together form a
mass just above this critical size. For a bomb explosion to occur, the
number of neutrons captured with fission must be greater than the
number of neutrons which escape. The number of neutrons which
escape depends on the surface area, whereas the number captured
depends upon the volume. As the quantity of fissionable material
increases, the volume increases faster than the surface area. Critical
size means the size at which the neutrons captured exceed the num-
ber which escape and fission occurs.
Plutonium is produced in nuclear reactors or "piles." The nuclear
reactor built at Oak Ridge, Tennessee, is essentially a large cube of
graphite bricks containing a number of horizontal channels into
which is placed pure uranium in the form of solid cylinders or
slugs enclosed in aluminum casings. Graphite is used to slow down
neutrons and is called a moderator. Heavy water is another good mod-
erator. Slow neutrons are more effective in producing fission than
are neutrons that travel at normal speeds. The bricks are built up
in layers, and since the structure was built by piling one layer of
bricks upon another, it is called anatomic pile.
A chain reaction is started with neutrons liberated from a bit of
beryllium mixed with radium in the center of the pile. The concen-
tration of neutrons is controlled by cadmium or boron-steel rods,
which absorb neutrons easily. Several nuclear reactions take place as
concrete shield
boron steel control rod
technician removing
tubes containing
radioactive
isotopes
protective lead shield
long graphite stringer
with
holes for aluminum tubes
Fig. 39. Simplified drawing of a
graphite-moderated atomic pile.
graphite moderator
aluminum
tubes
containing
uranium
servicing
elevator
second
floor
Adapted from a drawing of the Atomic Enerf/y Commission
shown on page 187. When the uranium slugs are ready for processing,
they are pushed out at the back of the reactor and new ones are fed in
at the front. The slugs fall into tanks of water where the U-239 grad-
ually changes into Np-239 and finally into plutonium. The slugs are
then dissolved in acid, and the plutonium is separated chemically
from the rest of the elements present. This chemical process of sepa-
ration is much easier than physical separation of U-235 from U-238.
Since dangerous radiations and radioactive material are produced
during these changes, all operations are performed by remote control.
The whole pile is surrounded with several feet of concrete to shield
and protect the operating personnel. Several reactors are in opera-
tion in this country. The one at Oak Ridge is air-cooled, while the
pile at Hanford, Washington, is water-cooled. Intensive research on
other coolants is now being carried on. Metals, such as sodium, in a
liquid state have been found useful for this purpose.
The atomic pile liberates tremendous amounts of heat. Efforts to
build industrial nuclear reactors which will utilize this huge source
casting
radioactive
cobalt
Co 60
Fig. 40. Use of radioactive cobalt for
the detection of flaws in castings.
defect in
casting
Adapted from a drawing of the
Atomic Energy Commission
film
developed film
shows defect
189
190
NEW WORLD OF CHEMISTRY
of power are already well under way. Atomic power plants "burning"
nuclear fuel will supply electricity, not only in our own country but
later on also, it is to be hoped, in those areas of the world which are
poor in coal, petroleum, and natural water power. Nuclear furnaces
may be built by private industry with uranium supplied by the
Atomic Energy Commission. The plutonium manufactured during
the process will be turned back to the United States Government.
Radioisotopes, first produced artificially by Madame Curie's
daughter. In 1934, Irene Joliot-Curie (zho'lyf/) and her husband,
Frederic Joliot-Curie, by bombarding boron with alpha particles,
produced a neutron and a radioactive isotope of nitrogen. Here was
another case of modern alchemy.
Boron + helium -
mass 10 + mass 4 -
> radioactive nitrogen -f neutron
» mass 13 + mass 1
Since then, scientists have made more than 700 new and different
radioisotopes in cyclotrons and nuclear reactors. For example, when
a bit of ordinary iodine (atomic weight 127) is placed in an atomic
pile where it is bombarded with neutrons, it changes to a radio-
isotope of iodine of mass 131. The Atomic Energy Commission sup-
plies hundreds of radioisotopes to research groups all over the world.
A new "tracer technique" uses radioisotopes. Research in medi-
cine, biology, agriculture, and many other fields has been helped tre-
mendously by this new method. Radioactive iodine, for example, is
being used in thyroid diagnosis and therapy. A person suffering from
hyperthyroidism is fed with a trace of sodium iodide containing I131.
With the aid of a Geiger counter, the rate at which this iodine com-
pound collects in the thyroid gland can be accurately determined.
Radioactive cobalt, Co"0, loses radioactivity in five days and is used
Atomic Ei
Materials may be made
radioactive by exposure
within a nuclear reactor
such as this water-boiler
reactor. Such radioisotopes
have many uses in indus-
try, agriculture, and medi-
cine.
NUCLEAR ENERGY AT LAST!
191
in cancer therapy as a substitute for radium and x-ray treatments.
Radioactive carbon, C1*, is a wonderful tool in the study of photo-
synthesis and such problems of human health- as sugar metabolism.
Radioactive phosphorus, P32, is used in agricultural research dealing
with the accumulation, utilization, and action of phosphate fertiliz-
ers. Industry is using radioisotopes in the improvement of steel, in
studying the action of catalysts, in measuring the flow of under-
ground water, oil, and gas, and in the detection of leaks.
Nuclear energy in the future? So far, nuclear energy has been
used mainly as a military weapon, and for research. No one knows
what the peaceful use of nuclear energy will bring. It seems likely
that nuclear science will give higher standards of living to all peoples.
The unlocking of almost unimaginable stores of energy should
teach man important lessons. Nuclear energy may transform the
world by improving the health arid raising the standards ot living of
millions of persons. But this same instrument in the form of a Hy-
drogen-bomb can destroy civilization as the A-bomb wiped out much
of Hiroshima and Nagasaki. Therefore, the nations of the world
must find a way of preventing a war with nuclear weapons from ever
taking place.
The H-bomb, based on the fusion of the heavy isotopes of hydro-
gen into helium and triggered by the 100,000,000°C. temperature of
an A-bomb explosion, could be made of unlimited size. This thermo-
nuclear reaction may be expressed as follows:
Deuterium + tritium — > helium + neutron -f- energy
H2 + H3 — > He +
n
The sudden conquest of nuclear energy demonstrated that science
in a democracy is strong and tremendously creative. By constant
vigilance, we must strive to keep it so.
Fig. 41. Simplified drawing of the use of nuclear energy for generating electricity.
Adapted from a drawing of the Atomic Energy Comnnsifion
Reactor control console
Electric power — «
ROWER PLANT
A. >*«,,. .,. Turbine
Reactor core
Uranium rods
NUCLEAR REACTOR
Oter
£tean
192
NEW WORLD OF CHEMISTRY
YOU WILL ENJOY READING
Curie, Eve. Madame Curie. Garden City Publishing Co.,
New York, 1943. In this intimate portrait of her mother, Eve
Curie tells an entrancing story of the discovery and isolation
of radium.
Glasstone, Samuel. Sourcebook on Atomic Energy. D. Van
Nostrand Company, New York, 1950. This book, written at
the request of the United States Atomic Energy Commission,
brings together in readable narrative the important facts about
the past history, recent status, and possible future of nuclear
energy.
Fermi, Laura. Atoms in the Family. The University of Chi-
cago Press, Chicago, 1954. A simple, intimate account of the
events surrounding the conquest of a controlled nuclear chain
reaction, written by the wife of Enrico Fermi.
Dean, Gordon. Report on the Atom. Alfred A. Knopf, New
York, 1953. An account of the role of nuclear energy by the
former head of the U. S. Atomic Energy Commission.
Smyth, Henry D. Atomic Energy for Military Purposes.
Princeton University Press, 1945. This is the so-called Smyth
Report released by the Army a few days after the first atomic
bomb was dropped on Japan. It is semitechnical in nature and
not easy reading.
USEFUL IDEAS DEVELOPED
1. X-rays consist of vibrations shorter than those of light.
They are produced by bombarding metals with a stream of
rapidly moving electrons. These x-rays have great penetrating
power.
2. Certain elements break down naturally, or disintegrate,
forming lighter elements including helium nuclei. This prop-
erty is called radioactivity.
3. Radium gives off alpha particles, beta particles, and
gamma rays, which are similar to x-rays.
4. The simplest units of matter thus far identified are the
electron, proton, neutron, positron, and meson.
5. The positron has the same mass as the electron; it has
a positive charge; and it can be formed during the disintegra-
tion of radioactive nitrogen.
6. Mesons have masses between the mass of an electron and
the mass of a proton; they have either a negative charge, a posi-
tive charge, or are neutral; they are produced by the bombard-
ment of atoms with cosmic rays or with rays of 100-million
electron-volts from the betatron.
NUCLEAR ENERGY AT LAST! 193
7. Modern alchemy, or artificial transmutation, is an ac-
complished fact.
8. Artificially radioactive elements were first obtained by
the Joliot-Curies by bombarding boron with alpha particles.
Hundreds of new radioisotopes have since been produced.
9. The age of nuclear energy was ushered in by the con-
trolled fission of U-235 and plutonium in 1945.
USING WHAT YOU HAVE LEARNED
Group A
1 . (a) Who discovered x-rays? (b) How are they produced?
(c) How do they differ from light?
2. How did Becquerel's discovery lead to the discovery of
Ra?
3. List some of the properties and uses of Ra.
4. What is meant by the half-life of an element?
5. What are the alpha, beta, and gamma rays emitted
during the disintegration of Ra?
6. Describe the construction and operation of a Geiger
counter.
7. Name five different particles that have been expelled
from the nucleus of atoms during bombardment.
8. Name several known facts concerning the mesons.
9. (a) Has the dream of transmutation come true? (b) Ex-
plain your answer.
10. Show by a diagram how Rutherford changed N into H.
11. (a) Write the Einstein equation for the conversion of
mass and energy, (b) Illustrate the meaning of this equation
in terms of the change of lithium into helium.
12. What is meant by a chain reaction?
13. How does artificial radioactivity differ from the natural
radioactivity of Ra?
14. Describe some of the events since 1938 leading up to the
final conquest of nuclear energy.
15. By means of three equations, explain the production of
plutonium from uranium.
16. By means of equations, explain how nuclear energy was
released in 1945.
194
NEW WORLD OF CHEMISTRY
17. Describe the construction of a nuclear pile.
18. By means of an equation, explain the explosion of a
hydrogen bomb.
Group B
19. "The conception of the structure of the atom makes it
possible for present-day scientists to explain the riddle of
transmutation." Explain this statement.
20. Make a diagram of the heaviest known element showing
the composition of the nucleus and the positions of all its
electrons. Consult your teacher or a recent edition of some col-
lege chemistry textbook.
21. Scientists believe that in releasing about four killowatt-
hours of energy in burning one pound of coal, a very small part
of the coal is converted into energy. Why has this fact not been
proved by experiment?
22. Describe briefly the future of the peacetime uses of
nuclear energy.
23. (a) Would you buy stock in a company organized to
exploit nuclear energy? (b) Give reasons for your answer.
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Make a small model of an electrical power plant utilizing
nuclear fuel. Consult your bibliography on page 192.
2. Visit your dentist or doctor and ask him to show you his
x-ray machine. Make a report on its construction and opera-
tion, using diagrams.
3. Consult your teacher of economics on the question:
"What would the effect of commercial transmutation of iron
into gold be on the financial structure of the world, if it were
accomplished tomorrow?" Write a report on this subject.
4. Compare the effects of the Industrial Revolution and the
possible effects of use of nuclear energy. Ask your social science
teacher for help.
5. Write a two-page report on the construction and opera-
tion of a cyclotron or other particle accelerator now in use.
13.
ACIDS:
HYDROCHLORIC ACID, A TYPICAL ACID
. . . To one man science is a sacred
goddess to whose service he is happy
to devote his life; to another she is a
cow who provides him with butter.
Liebig, 1803-1882
Gunfire in the American wilderness helped us to learn more
about digestion. In. 1822 at a remote fort on Mackinac Island be-
tween Lake Huron and Lake Michigan, a French-Canadian, Alexis St.
Martin, was brought in tor medical treatment. An accidentally dis-
charged musket had sent a bullet through the wall of his stomach.
Dr. William Beaumont, an army surgeon, patched him up.
Despite great effort, it was impossible to get the wound to close,
and on healing, a flap covering an opening into St. Martin's stomach
was left. Through this opening Beaumont could reach directly into
St. Martin's stomach. Beaumont got a strange idea. This freak "lid"
over the hole into St. Martin's stomach would enable him to perform
experiments to discover the digestive action of the juices of the
stomach. St. Martin was agreeable and Beaumont tied pieces of food
to a string, inserted them into St. Martin's stomach and, after several
hours, removed what was left of the food.
In this way Beaumont gave science the first accurate facts concern-
ing the relative digestibility of foods and the composition of gastric
juice. He found gastric juice to contain a small amount of hydro-
chloric acid (about 0.3 percent) , which helps to digest certain foods,
especially proteins.
195
196
NEW WORLD OF CHEMISTRY
Fig. 42. Laboratory prepara-
tion of hydrogen chloride. Why
is the end of the delivery tube
above rather than below the
level of the water?
x water
How hydrogen chloride is prepared in the laboratory. To prepare
hydrogen chloride in the laboratory, concentrated sulfuric acid is
added to sodium chloride in a flask, as shown in the illustration
above. This mixture is heated, and hydrogen chloride gas, which
is liberated readily, is collected by the displacement of air. In this
chemical change, a double replacement occurs, as indicated in the
following equation:
H2SO4 + NaCl -» HC1 + NaHSO4 (sodium bisulfate)
At the outset, we must distinguish clearly between HC1 (hydro-
gen chloride gas) and HC1 (hydrochloric acid) . When hydrogen
chloride gas is dissolved in water, hydrochloric acid is obtained.
Both are represented by the same formula, but their physical and
chemical properties are entirely different.
How hydrogen chloride is prepared commercially. When a jet of
hydrogen is burned in chlorine, hydrogen chloride gas is formed (see
Chapter 10) .
H2 + C12 -> 2HC1
This is one commercial method of manufacturing HC1. A second
commercial method is similar to the laboratory one but at a some-
what higher temperature producing Na2SO4 instead of NaHSO4.
2NaCl + H2SO4 -> 2HC1 + Na2SO4
Chemical properties of hydrochloric acid. Hydrochloric acid is
one of the most common and useful laboratory chemicals, or re-
agents. Some of its important properties are:
1) Taste. Dilute hydrochloric acid has a sour taste.
2) Action on indicators. Hydrochloric acid reacts with a group of
substances known as indicators, causing a color change. For example,
hydrochloric acid turns blue litmus pink. It also turns reddish-purple
phenolphthalein (fe-nol-thal'en) colorless.
ACIDS: HYDROCHLORIC ACID 197
3) Action with metals. When hydrochloric acid is in contact with
most metals, a reaction takes place. Hydrogen is liberated (recall the
laboratory method for the preparation of hydrogen) , and chlorides
are formed. Note that the metal replaces the hydrogen of the HC1.
2HC1 + Zn -> H2 t + ZnCl2 (zinc chloride)
2HC1 + Fe -> H2 f + FeCl2 (ferrous chloride)
4) Action with bases, or hydroxides of metals. Hydrochloric acid
reacts with a base, forming a neutral compound that possesses the
qualities of neither acid nor base (bases are discussed in Chapter 14) .
Pure water is the only other product of this reaction. For example,
when hydrochloric acid and sodium hydroxide react, the products
are common salt and water.
HC1 + NaOH -» NaCl + H2O
Why hydrochloric acid is a typical acid. The chemical proper-
ties of hydrochloric acid are characteristic of the whole group of com-
pounds known as acids. We may now define an acid as a water solu-
tion of a compound with the following characteristic properties:
1) An acid has a sour taste. The sour taste, or tartness, of fruits is
caused by certain acids, such as citric acid, which is found in lemons,
limes, and grapefruit.
2) An acid turns blue litmus pink, reddish-purple phenolphthalein
colorless, and acts with other indicators in the same way as hydro-
chloric acid.
3) An acid contains hydrogen that can be replaced by most metals,
forming compounds known as salts. This does not mean that all com-
pounds containing hydrogen are acids. Sugar, for example, contains
hydrogen, but it is not an acid, because its hydrogen cannot be re-
placed by metals. A compound in which the hydrogen of an acid
has been replaced by a metal is known as the salt of an acid. Thus
sodium chloride, NaCl, is a salt of hydrochloric acid, and sodium ni-
trate, NaNO3, is a salt of nitric acid, HNO3.
4) An acid neutralizes any base, forming water and a salt whose
composition depends on both the acid and the base used.
In general, an acid consists of hydrogen and a nonmetallic ele-
ment, or hydrogen and a radical. The hydrogen of the acid can be
replaced by a metal. Strictly speaking, while certain substances show
acid properties only in water solution, they are commonly called
acids even when dry. Thus, perfectly dry H2SO4 is called sulfuric
acid, even though it exhibits acid properties only when it is in a
Corn Industries Research Foundation, Inc.
In these converters, corn starch is changed, under heat and pressure, into corn
sugar, or dextrose. Hydrochloric acid serves as a catalyst in the process. Because
the process simulates human digestion, dextrose is, in effect, a predigested food
and is readily assimilated by the body.
water solution. Bases also exhibit their characteristic basic properties
only when water is present. Certain characteristics of acids or bases
in solution are discussed in Chapter 16.
The general method for the preparation of an acid. Sulfuric acid
is one of the most important acids, because it is used as a raw material
in the manufacture of most of the other acids. Sulfuric acid has two
special characteristics that make it well suited for this purpose:
(1) its low cost, and (2) its high boiling point (338°C.) . The manu-
facture of sulfuric acid is discussed on page 309.
The preparation of hydrochloric acid illustrates the general
method used in preparing acids from sulfuric acid. First, a salt of
the acid to be prepared is chosen as a source of the nonmetallic ele-
ment of the acid. Common salt is the least expensive and most abun-
dant source of chlorine. Many other chlorides can be used, but they
are more expensive. Sulfuric acid supplies the hydrogen.
The salt of the acid to be prepared, NaCl, and sulfuric acid are
heated together gently. Hydrogen chloride gas is produced by the
reaction and is driven off and dissolved in water in the receiving
vessel, forming the acid. The higher boiling points of the other com-
pounds taking part in the reaction prevent their vaporizing and thus
keep them from passing over into the receiving vessel.
Many other acids are manufactured by treating their least expen-
sive and most abundant salts with sulfuric acid. The acid formed is
198
ACIDS: HYDROCHLORIC ACID
199
usually separated from the reacting substances by methods based on
the differences in their boiling points.
Physical properties of hydrochloric acid. Hydrochloric acid is a
colorless liquid, heavier than water. That is, its specific gravity is
greater than one. It possesses an irritating odor. Both the boiling-
point and the specific gravity of hydrochloric acid are determined by
the weight of hydrogen chloride gas dissolved in the water. Hydro-
chloric acid containing 20 percent hydrogen chloride gas by weight
boils at 110°C. Impure hydrochloric acid is called muriatic acid and
is usually yellow in color. In this form it was known for many years
before Priestley, in 1772, first isolated pure hydrochloric acid.
Hydrochloric acid cleans metals. Before coating metals, such as
iron and steel, with plates, films, or coatings of other metals, includ-
ing chromium, silver, tin, and zinc, the surface of the metal must be
clean and free of oxides. Removing oxides and otherwise cleaning
the surface of a metal to be plated or coated is a process known com-
mercially as pickling. One of the chief industrial uses of hydro-
chloric acid is in the pickling of metals, especially before coating
with tin in tinning, or xinc in galvanizing, or with the materials that,
after firing, result in enamelware.
Small quantities of hydrochloric acid, usually muriatic acid, are
used in removing rust stains from vitreous washbasins and lavatories.
Plumbers often use muriatic acid as a flux before soldering.
Racks of sheet steel emerge from the pickler. The worker in the foreground is
dipping the sheets in water to remove the acid.
200 NEW WORLD OF CHEMISTRY
Hydrochloric acid is used in making other chemicals. Chlorides
of many metals, including magnesium chloride, aluminum chloride,
and zinc chloride, are made by the reaction of hydrochloric acid and
a carbonate or oxide of the metal. For the most part, such chemicals
are of very high quality, and are used chiefly by manufacturing chem-
ists and drug houses, and by druggists. Zinc chloride is used to im-
pregnate wood to prevent decay, in soldering, and in flame-proofing.
Silver chloride, AgCl, one of the several light-sensitive silver com-
pounds used to coat photographic film may be made by the reaction
of silver nitrate, AgNO3, and hydrochloric acid. In the manufacture
of glucose from starch, hydrochloric acid acts as a catalytic agent. It
is used also in large quantities in the manufacture of glue and gela-
tin, in the purification of boneblack, and in the processing of textiles.
Physical properties of hydrogen chloride gas. Hydrogen chloride
gas is colorless, heavier than air, and has a sharp, penetrating odor.
It is extremely soluble in water. If a test tube of hydrogen chloride
gas is placed mouth downward in water, the water will rise almost
to the top, as the gas dissolves. Water dissolves about 500 times its
own volume of this gas under normal laboratory conditions. The
gas can be liquefied and solidified, just as all other gases.
Chemical properties of hydrogen chloride gas. Hydrogen chloride
does not show acid properties unless it is dissolved in water or unless
some water vapor is present. When dry, it is completely inactive. Its
attraction for water is so great that it forms a cloud or mist in moist
air of tiny droplets of hydrogen chloride solution.
This property may be used as a test for hydrogen chloride gas. If
you blow across the mouth of a bottle containing concentrated hy-
drochloric acid, a mist will form. This mist is caused by the moisture
in your breath combining with the hydrogen chloride vapor that
rises from the bottle. Hydrogen chloride does not burn.
Preparation and properties of hydrofluoric acid, HF. The prep-
aration of hydrofluoric acid follows the general method for making
an acid. Sulfuric acid reacts with calcium fluoride, CaFo, the most
common salt of hydrofluoric acid, forming hydrofluoric acid. This is
the method used by Scheele when he first prepared it in 1771. Cal-
cium fluoride is the chief constituent of the mineral fluorspar, found
in several parts of the United States.
H2SO4 + CaF2 -> 2HF | + CaSO4
Pure anhydrous hydrogen fluoride is a colorless gas which boils at
room temperature (19.5°C.) . It is deadly if inhaled. It dissolves
in water, forming a colorless acid that vaporizes at low temperatures.
Calibration marks are carefully
scratched through the wax coat-
ing on a graduated glass cylinder
prior to exposing it to hydrofluoric
acid.
Corning (Hass Works
Such an acid is called a fuming acid. Since hydrofluoric acid re-
acts with glass, quartz, and most metals, it is distilled in dishes made
of lead and must he kept in bottles made of polyethylene or other
plastics with which it does not react. This acid produces powerful
burns by poisoning the tissues.
How glass is etched with hydrofluoric acid. Etching is the biting
out of particles of glass or metal by means of chemicals. Hydrofluoric
acid etches glass because it unites with the silicon dioxide, SiO2, of
the glass, forming silicon tetrafluoride, SiF.,, which is a gas.
SiO2 + 4HF -> SiF4 1 + 2H2O
In etching glass articles, such as thermometers, electric-light bulbs,
and windows, the surface is completely covered with wax and the
design to be etched is scratched through the wax. The object is then
brought into contact with the vapor of hydrofluoric acid. When the
action on the exposed glass has gone as far as necessary, the object is
removed from contact with the vapor. In frosting the inside of an
electric-light bulb, a small amount of hydrofluoric acid is poured
into the bulb, shaken for an instant, and poured out, and the bulb
is thoroughly rinsed. It is also used in making Freon refrigerants, and
as a catalyst in the manufacture of high octane gasoline.
The other halogen acids. Theoretically, both hydrobromic acid.
HBr, and hydriodic acid, HI, may be prepared by the general method
used in preparing acids.
H2SO4 + KBr -» HBr | + KHSO4
H2SO4 + Nal -> HI j + NaHSO4
201
202
NEW WORLD OF CHEMISTRY
When first formed they are colorless gases with strong irritating
odors, but are almost immediately oxidized by the oxygen of the air,
forming free bromine and free iodine. In water solution, hydro-
bromic and hydriodic acids are unstable; on exposure to air they
decompose as a result of oxidation.
4HBr + O2 -» 2H2O + 2Br2
Heat of formation and chemical stability. When hydrogen and
chlorine react, forming hydrogen chloride, a great deal of heat is
produced. A reaction in which heat is liberated is called an exother-
mic reaction (ex — out, therme — heat) . Exothermic reactions con-
tinue after they are first started by external heat. On the other hand,
when hydrogen and iodine unite, heat is continuously absorbed and
heat must be added if the reaction is to continue. A reaction in
which heat is absorbed is an endothermic reaction.
The number of calories of heat absorbed or liberated, during the
formation of a mole (see page 130) of an element or compound,
is called its heat of formation. A substance which liberates heat dur-
ing its formation is said to have a positive heat of formation; a sub-
stance which absorbs heat is said to have a negative heat of forma-
tion. A compound such as sodium chloride, NaCl, whose positive
heat of formation is very great (97,800 calories) iswery stable. Hy-
drobromic acid, HBr, whose positive heat of formation is small (8400
calories) is somewhat unstable. Hydriodic acid, HI, which has a
negative heat of formation, is very unstable.
A knowledge of the heats of formation is very useful to chemists.
For example, we can tell whether a certain compound will form and
how easily it can be prepared. A compound formed by replacement
has a higher positive heat of formation than the compound from
which it is formed and hence is more stable.
2KBr
C12
2KC1
Br2
By similar reasoning we can see why bromine will not replace chlo-
rine from KC1.
ACIDS: HYDROCHLORIC ACID
203
YOU WILL ENJOY READING
Clendening, Logan. The Human Body* pp. 74-75. Alfred A.
Knopf, New York, revised edition. 1945. An amusing story of
Dr. William Beaumont's experiments with Alexis St. Martin.
Jaffe, Bernard. Men of Science in America, pp. 157-158.
Simon and Schuster, New York, 1944. The dramatic story of
the pioneer investigations of the American doctor, William
Beaumont, and of John R. Young, which resulted in our
knowledge of the presence of hydrochloric acid in the gastric
juice.
Lowry, T. M. Historical Introduction to Chemistry,
pp. 12-16. The Macmillan Co., London, 1915. The discovery
of the common acids and their chemical and physical proper-
ties are reviewed.
USEFUL IDEAS DEVELOPED
1. An acid is a substance whose water solution (1) has a
sour taste, (2) turns blue litmus pink, (3) contains hydrogen
that can be replaced by each of many metals with the forma-
tion of a salt and the liberation of hydrogen, and (4) neu-
tralizes bases, forming a salt and water.
2. Chemical reactions in which heat is liberated are called
exothermic reactions; those in which heat is absorbed are
called endothermic reactions.
3. The heat of formation of a compound is the number of
calories of heat liberated or absorbed in the formation of
1 gram-molecular weight of the compound.
4. Compounds with high positive heats of formation are
stable; compounds with low positive heats of formation are un-
stable; compounds with negative heats of formation are very
unstable.
USING WHAT YOU HAVE LEARNED
Group A
1. Describe the laboratory method of preparing HC1.
2. (a) What are the two methods of preparing hydrogen
chloride commercially? Write an equation for each method.
(b) What type of chemical reaction does each method repre-
sent?
204 NEW WORLD OF CHEMISTRY
3. Give the properties of HC1 as to (a) taste, (b) action on
indicators, (c) action on metals, (d) action on oxides, (e) ac-
tion on hydroxides of metals.
4. (a) What is a salt? (b) Give three examples.
5. (a) What are the salts of HC1 called? (b) Name three.
6. (a) What are salts of HNO3 called? (b) Name one.
7. (a) What is an acid? (b) Name five acids.
8. Make a list of the properties of hydrogen chloride gas.
9. What is the percentage of hydrogen chloride in a solu-
tion that has a fixed boiling point of 110°C.?
10. What weight of pure salt is needed to prepare 292 Ib. of
hydrochloric acid containing 15 percent HC1 by weight?
1 1 . Compare the solubility of hydrogen chloride in H2O
with the solubility of air, O2, H2, CO2, and N2.
12. What causes the fuming of hydrogen chloride in moist
air?
13. What test or tests would you use in showing the pres-
ence of HC1 in a liquid?
14. (a) What are the uses of HG1? (b) What process in the
human body depends in part upon HC1?
15. (a) State the general method of preparing acids.
(b) Why is H2SO4 so generally used?
16. Scheele was the first to prepare HF. His method is still
used today. Describe it.
17. By means of an equation, explain the etching action
of HF.
18. Explain how HF is stored.
19. Compare the properties of the other halogen acids with
those of HC1.
20. (a) What is a calory? (b) What is heat of formation?
21. (a) What is the difference between an exothermic re-
action and an endothermic reaction? (b) Give an example of
each.
22. Explain the relationship between the heat of formation
and the stability of a substance.
23. The positive heat of formation of H2S is 2730 calories.
(a) Is H.,S a stable compound? (b) Explain.
ACIDS: HYDROCHLORIC ACID 205
24. Write balanced equations for the following:
a) Copper oxide + hydrochloric acid — *
b) Hydriodic acid 4- oxygen — >
c) Sulfuric acid + sodium bromide — »
Group B
25. What is the effect of chlorine water on litmus?
26. When HC1 is boiled, what passes off? Explain.
27. A bottle of HI turns brown. Explain.
28. What chemical tests would you use in identifying each
of the halogen acids?
29. Why is muriatic acid used in the soldering of metals?
30. Water may be considered an acid. Explain.
31. Why is the general method for preparing acids not used
in the preparation of HI?
32. In the electrolysis of HC1, the volumes of C12 and H2
collected are not the same. Explain.
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Take home some strips of litmus paper and test various
foods and other substances found at home for the presence of
acids. Report your findings to the class. Make a list of common
foods that react to your test as acids; as bases. What do you
conclude from your tests?
2. Prepare a report on the experiments performed by Beau-
mont with Alexis St. Martin which dealt with the discovery of
HC1 in gastric juice.
3. Consider carefully the quotation at the beginning of this
chapter. Prepare a report for the class on the meaning of this
quotation. Illustrate with examples from the lives of great
scientists and inventors.
14.
BASES:
SODIUM HYDROXIDE, A TYPICAL BASE
. . . Rouelle [the teacher of Lavoi-
sier], to whom we owe the term base,
described how natural salts had been
restricted at first to salts formed by
the union of acids with alkalies
which are soluble in water, and im-
part on the tongue a saline taste.
T/M. Lowry, 1915
How sodium hydroxide is prepared in the laboratory. The lab-
oratory preparation of hydrogen by the action of sodium on water
was discussed in Chapter 3. Each sodium atom replaces one of the
atoms of hydrogen in the water molecule.
2Na + 2HOH -> SNaOH + H2 1
After the water evaporates, a white solid, sodium hydroxide, is left.
Chief properties of sodium hydroxide. Sodium hydroxide is a
common and useful substance. Some of its important properties are:
1) Action on indicators. Sodium hydroxide turns pink litmus blue
and colorless phenolphthalein reddish-purple.
2) Action with acids. Sodium hydroxide reacts with an acid, form-
ing a salt, thus causing the original properties of the acid to disap-
pear along with its own. For example, sodium hydroxide unites with
hydrochloric acid, forming common salt and water. The presence of
the salt may be proved by evaporating the water and tasting the
solid left behind.
3) Feel. A water solution of sodium hydroxide has a slippery,
soapy feeling.
206
BASES: SODIUM HYDROXIDE 207
Why sodium hydroxide is a typical base. The chief properties of
sodium hydroxide are characteristic of the whole group of com-
pounds known as bases. We may now define a base as a compound
(1) that contains the hydroxyl group (OH) , (2) whose water solu-
tion is soapy to the touch, (3) whose water solution turns pink
litmus blue and colorless phenolphthalein reddish-purple, and (4)
whose water solution reacts with acids, forming water and a salt of
the acid used.
In general, then, a base consists of a metallic element or radical and
one or more hydroxyl groups (OH) .
Strong bases, such as sodium hydroxide and potassium hydroxide,
are often known as alkalies. Three other important bases are cal-
cium hydroxide (slaked lime) , ammonium hydroxide (ammonia
water) , and magnesium hydroxide (milk of magnesia) .
How bases react with acids. When a base reacts with an acid, the
hydrogen of the acid combines with the hydroxyl radical of the
base, forming water, and the metal or metallic radical of the base
unites with the nonmetal or acid radical of the acid, forming a salt.
Such a chemical change is called neutralization. All neutralization
reactions are double replacements.
HC1 + NaOH -» HOH + NaCl (a salt)
HNOa + KOH -* HOH + KNO3 (a salt)
H2SO4 + Ca(OH)2 -» 2HOH + CaSO4 (a salt)
HBr + NH4OH -> HOH + NH4Br (a salt)
Neutralization is the reaction of an acid with a base, forming wa-
ter and a salt. A salt, then, is a compound made up of a metal or a
metallic radical and a nonmetal or an acid radical.
Titration and the use of molar and normal solutions. It is some-
times necessary to know how much of an acid, a base, or a salt is
present in a solution. For example, we may want to know how much
acetic acid is present in a given sample of vinegar. One way is to
neutralize a given volume of this liquid with a solution of a base
whose composition is known. This process of determining the
strength of an acid or a base with the help of neutralization reaction
is called titration. It is carried out in long tubes (burettes) which
are measured off in milliliters so that volumes can be read off di-
rectly. A definite volume of the acid solution to be tested is neutral-
ized with a standard solution of a base, that is, one whose composi-
tion is known. The point at which neutralization is reached (end
point) is determined by the use of an indicator such as phenolphtha-
lein which suddenly changes color.
NEW WORLD OF CHEMISTRY
liquid
Fig. 43. Titration. The degree of acidity of
any acid solution is determined by measur-
ing the amount of normal basic solution
required to neutralize it. The process is
reversed to find the strength of any basic
solution.
-* — burette
beaker
tT *%&(*
„* -"
There are two kinds of standard solutions used, namely, molar
solutions and normal solutions. One liter of a molar solution con-
tains one mole or gram-molecular weight of dissolved substance. For
example, a molar solution of Nad contains 23 + 35.5, or 58.5, grams
of NaCl in one liter of this solution. A molar solution of Ca (OH) 2
contains 74 g.
One liter of a normal acid solution contains one gram of replace-
able hydrogen. One liter of a normal basic solution contains 17 grams
of the OH group, and one liter of a normal salt solution contains the
equivalent of one gram of replaceable hydrogen. For example, one
liter of a normal solution of H2SO4 contains 98 -4- 2, or 49, grams
H2SO4 since this acid contains two replaceable hydrogen atoms. One
liter of a normal Ca (OH) 2 solution contains 74 -:- 2 grams Ca (OH) 2.
One liter of a normal A1C18 solution contains 133.5 -4- 3 grams A1C1S
since this salt contains three hydrogen equivalents. Decimal frac-
tions are used in referring to both molar (M) and normal (N) solu-
tions. Thus a 0.1M solution of HC1 contains 36.5 -~ 10, or 3.65,
grams HC1 per liter of solution.
PRACTICE WORK ON MOLAR AND NORMAL SOLUTIONS
1. How much NaCl per liter of solution does a 0.3M solu-
tion of NaCl contain?
2. How much acid, base, or salt per liter of solution do the
following solutions contain: 0.1N AgNO3, 0.3N MgBr2,
1.5N H3PO4, 0.125N CuCl2?
3. How would you prepare the following solutions:
0.8N HNO3, 3.0N KOH, 0.25N Mg(OH),
1.3N A1(OH)3?
BASES; SODIUM HYDROXIDE 209
Heat of neutralization. All neutralization reactions are exothermic.
When any strong acid and strong base react, a definite quantity of
heat is liberated, 13,700 calories per gram-molecular weight, or mole,
of water formed. This heat is called heat of neutralization. When
36.5 grams of hydrochloric acid react with 40 grams of sodium hy-
droxide, 18 grams of water and 58.5 grams of sodium chloride are
formed with the liberation of 13,700 calories. In the case of the ac-
tion of sulfuric acid on potassium hydroxide, 2 X 13,700 calories are
produced, because two moles of water are formed,
H2SO4 + 2KOH -» K2SO4 + 2H2O
Important uses of neutralization. After petroleum has been puri-
fied with the aid of sulfuric acid, the excess acid is removed by neu-
tralizing it with a base, usually sodium hydroxide. Much of the
soil in the United States is sour (acid) . Four-fifths of the cultivated
land in the central western states is sour and therefore not fully
productive. These acid soils may be neutralized by the addition of
lime, CaO, which combines with water, forming a base. Calcium car-
bonate, CaCOa, is used also. Soil with excess lime is neutralized with
ammonium sulfate, (NH4) 2SO4.
The destructive effect of acid stains or burns may be minimized
by the prompt application of either a weak base or sodium bicar-
bonate, which neutralizes the effect of the acid. Similarly, an alkali
burn is treated by application of a mild acid, such as boric acid or
vinegar. An excessive acid mouth condition is sometimes treated by
using small amounts of milk of magnesia, Mg (OH) 2, a mild base, or
antacid.
Normal blood is slightly alkaline. Alkaline reserve refers to the
amount of base present as bicarbonate in the blood. Even slight
changes in the normal alkalinity of the blood result in serious body
disturbances. Such disturbances may, in some cases, be corrected by
using neutralizing chemicals administered by a physician.
Existence of an alkaline reserve (a scientific fact) is, unfortunately,
used by some advertisers as a reason for selling to the public huge
quantities of "alkalizers" to dose ailments that are best treated by
other methods administered by a competent physician. This is an
excellent example of the abuse of scientific knowledge, against which
intelligent persons must always be on guard.
Physical properties of sodium hydroxide. The white solid, sodium
hydroxide, is very soluble in water. On exposure to moist air, it ab-
sorbs large quantities of water and changes into a pool of sodium
hydroxide solution. This property of deliquescence makes sodium
hydroxide useful as a drying agent. Usually it is sold in pellets that
must be kept in well-sealed bottles which exclude moisture.
Chemical properties of sodium hydroxide. In addition to typical
properties of bases already mentioned, sodium hydroxide has other
chemical properties. When exposed to air, it unites with carbon di-
oxide, forming sodium carbonate.
2NaOH + CO2 -> Na2CO3 + H2O
The common name for sodium hydroxide is caustic soda, given to
it because of its caustic, or burning, action upon the skin. It dis-
solves wool but has little effect upon cotton, rayons, or nylons. Be-
cause of this fact, it is used in determining the amount of wool in a
cotton-wool mixture.
Potassium hydroxide (caustic potash} closely resembles sodium hy-
droxide in both chemical and physical properties.
Sodium hydroxide helps clothe us. Sodium hydroxide is one of
the most useful compounds known. It serves man in many ways,
chief among which is in the manufacture of rayon and cellulose
films similar to cellophane. Approximately one-fifth of the 3 million
tons of sodium hydroxide produced in a recent normal year in this
country was consumed in this way. The chemistry of the production
of rayon and cellulose films is discussed in detail in Chapter 37.
Many fabrics are made of cotton that has been treated with so-
dium hydroxide. When cotton fibers are placed in a solution of
sodium hydroxide, they lose part of their natural twist and acquire
a gloss that is considered very desirable by many persons. Cotton so
treated is known as mercerized cotton.
Sodium hydroxide helps keep us clean. Until very recently sodium
hydroxide and potassium hydroxide were used in making almost all
the cleansing agents for both industrial and home use. When sodium
210
hydroxide or potassium hydroxide reacts with a fat, soap and glycerin
are formed. Each year the soap industry uses thousands of tons of
sodium hydroxide and considerable amounts of potassium hydroxide
Another important use of sodium hydroxide is in the form of he
which contains about 94 percent sodium hydroxide. It is a useful
household cleansing agent because it dissolves grease. Lye produces
pamlul burns if it comes in contact with the skin. It should be used
cautiously and stored out of the reach of small children
When kitchen or bathroom plumbing becomes clogged, a strong
solution of lye is sometimes poured down the pipes. In this way
greasy, fatty accumulations are saponified and become soluble in wa-
ter. "Drano" and other common plumbing cleaners sold under vari-
ous trademarks contain percentages of impure sodium hydroxide
Sodium hydroxide helps peel fruits and vegetables. Many fruits
and vegetables are peeled before canning or dehydration by the use
of sodium hydroxide, or lye. Most of the large lye peelers are of the
moving conveyor type, and the fruits or vegetables are immersed in a
20- to 25-percent lye solution for from two to five minutes. Durino
this time, the lye solution reacts with the pectins, or binding mate*
rials, between the individual cells. The skins become loose and may
be removed by washing which also removes all traces of lye.
This process is an outgrowth of the making of old-fashioned lye
hominy, a staple in the diet of American pioneers. In the making of
lye hominy, grains of corn are soaked in a lye solution until the hard,
tough skin of the corn grain becomes loosened. Washing with fresh
water removes the skin and also the lye.
Other uses of sodium hydroxide. Large quantities of sodium hy-
droxide are used in reclaiming aluminum and salvaged rubber, in
the processing of many vegetable oils, and in the production of gas-
oline. In the refining of petroleum, large quantities of this "basic
211
212
NEW WORLD OF CHEMISTRY
heavy chemical" are used to neutralize the sulfuric acid with which
petroleum is purified. Of only slightly less importance is the use of
sodium hydroxide in digesting and purifying the cellulose of wood
pulp that is used in manufacturing paper. Potassium hydroxide is
the electrolyte in certain types of storage batteries such as the Edison,
and Ni-Cd batteries. Lithium hydroxide is used in submarines to
absorb CO2.
How sodium hydroxide is prepared for industrial use. Chlorine,
hydrogen, and sodium hydroxide are all formed at the same time dur-
ing the electrolysis of brine (see Chapter 10) . Both chlorine and so-
dium hydroxide are prepared industrially by this method. If chlo-
rine is the chief product desired, then sodium hydroxide is the
byproduct. That is, it is a substance formed incidentally during the
preparation of another substance. If sodium hydroxide is the com-
pound being manufactured, then chlorine is the byproduct.
The apparatus that is commonly used in the industrial prepara-
tion of sodium hydroxide by the electrolytic process is the Hooker
cell. The Nelson diaphragm cell, Vorce cell, and mercury cell are
also used. The graphite anodes (see Fig. 44) are covered by a so-
dium chloride solution and are separated from the cathode by an
asbestos diaphragm, which prevents the chlorine from mixing with
the sodium hydroxide. Chlorine gas escapes through an outlet in the
dome at the top of the cell, and hydrogen gas passes through an out-
let from the steel screen cathode. Because the sodium hydroxide
solution is heavier than the salt solution, it concentrates with it at
the bottom of the cell. It is drawn off, and evaporated to dryness
during which almost all of the NaCl crystallizes out. This process
Rased on a diagram by
Hooker Electrochemical Company
hydrogen
chlorine
brine
graphite anode
sodii
hydrc
Fig. 44. Hooker cell used in the pro-
duction of sodium hydroxide and
chlorine. Hydrogen is a byproduct.
BASES: SODIUM HYDROXIDE 213
is continuous, and more brine is added as the strength of the sodium
chloride solution diminishes.
2NaCl + 2H2O -» 2NaOH + H2 1 + C12 1
An older method, still very widely used, depends upon the conver-
sion of the cheaper base, calcium hydroxide, into sodium hydroxide
by means of a solution of sodium carbonate. Until 1940, more so-
dium hydroxide was produced by this process than by electrolysis.
Ca(OH)2 + Na2CO3 -» CaCO3 [ + 2NaOH
Since calcium carbonate is insoluble, it is separated from the soluble
sodium hydroxide by filtration.
Methods of preparing a salt. Most inorganic compounds, such as
sodium chloride, sodium nitrate, copper sulfate, and so forth, are
salts. We have already had occasion to refer to six of the seven meth-
ods of preparing salts. A list of these seven methods follows:
1) Neutralization:
HC1 + NaOH -» H2O + NaCl
2) Action of an acid on a metal:
3) Union of a metal and a nonmetal:
Fe-f S-+FeS
4) Action of an acid on the oxide of a metal:
2HC1 + CuO -> H2O + CuCl2
5) Action of an acid on a salt of a more volatile acid:
H2SO4 + 2NaCl -> 2HC1 + Na2SO4
6) Action of one salt on another salt:
AgNO3 + NaCl -> AgCl + NaNO3
7) Action of the oxide of a metal (basic oxide) on the oxide of
a nonmetal (acidic oxide) :
CaO + SO8 -» CaSO4
The horizontal lines under certain formulas indicate salts.
214 NEW WORLD OF CHEMISTRY
YOU WILL ENJOY READING
Fabre, Jean H. The Wonder Book of Chemistry, pp. 154-
170. Albert & Charles Boni, New York, 1922. Salts and neu-
tralization are discussed in a captivating manner.
Jaffe, Bernard. Chemical Calculations, pp. 96-102. World
Book Co., Yonkers, New York, 1947. Normal and molar solu-
tions and problems involving them are included.
USEFUL IDEAS DEVELOPED
1. A base is a substance that contains a metal or metallic
radical and one or more hydroxyl groups. Its water solution
is soapy to the touch, turns pink litmus blue, and reacts with
acids, forming water and a salt.
2. Neutralization is the action of an acid with a base, form-
ing water and a salt. The hydrogen of the acid unites with the
hydroxyl radical of the base, forming water.
3. When a strong acid and a strong base react, forming a
mole of water, 13,700 calories of heat are liberated. This
amount of heat is called heat of neutralization.
4. A salt is a compound made up of a metal or a metallic
radical and a nonmetal or an acid radical.
5. The seven methods of. preparing salts ar^: (1) neutraliza-
tion, (2) action of an acid on a metal, (3) union of a metal
and a nonmetal, (4) action of an acid on a metallic oxide,
(5) action of an acid on a salt of a more volatile acid, (6) re-
action between two salts, and (7) reaction of a basic oxide on
an acidic oxide.
USING WHAT YOU HAVE LEARNED
Group A
1. (a) Write an equation illustrating the method of pre-
paring NaOH in the laboratory, (b) How can the NaOH
formed be changed into a solid?
2. (a) What other base can be made in the same way?
(b) Write the equation for the preparation of this other base.
(c) What is its common name?
3. (a) What is a base? (b) List the four properties by
which a base can be recognized.
4. Give the names and formulas of three bases other than
NaOH that are often found at home.
5. (a) What is a salt? (b) What is neutralization? (c) How
is mercerized cotton prepared?
BASES: SODIUM HYDROXIDE
215
6. A soil is found to be acid. To obtain the best results
with certain crops a neutral soil is required. How would a
farmer correct the acid condition of his soil?
0.5M AlBr3, and
7. Write four equations illustrating neutralization.
8. (a) What is a molar solution? (b) What weight of acid,
base, or salt do the following solutions contain:
(1) 3M NaOH, (2) 0.1M H2SO4 (3)
(4) 1.8M HN03?
9. (a) What is a normal solution? (b) What weight of
acid, base, or salt do the following solutions contain:
(1) 0.1N Mg(OH)2, (2) 0.5N H,PO4, (3) 2.5N AgNO3,
and (4) 1.5N A1C13?
10. What is meant by titration?
T
11. Describe the electrolytic process for making NaOH,
illustrating your description with a diagram.
12. In a process of the kind mentioned in exercise 11, what
determines which is the product and which the byproduct?
13. Write the equation for a method of preparing sodium
hydroxide other than the electrolytic method.
14. (a) What is the annual consumption of NaOH in the
United States? (b) What are the three chief uses of NaOH?
15. (a) What is lye? (b) Why is it useful in the home?
(c) What cautions should be exercised in its use?
16. A sample of fabric was placed in a test tube containing
a solution of NaOH, and boiled for a few minutes. Half of the
fabric disappeared. What can you say about its composition?
1 7. Copy and complete the following table, inserting the cor-
rect formulas for the salts formed. Do not write in this book.
HC1
HNO,
H2SO4
H2C08
NaOH
NaCl
Ca(OH)2
NH4OH
(NH4)2C08
18. Classify as acids, bases, or salts: A1(OH)8, K2CO8
CuSO,, Pb (OH) „ HC2H802, CaCl2, BaSO4, HCN, H2S.
216 NEW WORLD OF CHEMISTRY
19. Which will produce a greater weight of alkali, 45 g. of
Na or 50 g. of K?
20. A piece of Na is thrown in H2O. The NaOH formed
exactly neutralizes 24.5 g. of H2SO4. What weight of Na was
used?
Group B
21. (a) Write the equations for the neutralization of sodium
hydroxide with (1) hydrochloric acid, and (2) sulfuric acid,
(b) The heat of neutralization for the first reaction is 13,700
calories; for the second reaction, 2 X 13,700 calories. Inspect
the two equations. Explain why the heat of neutralization is
twice as large in the second reaction.
22. What volume of a 0.1N KOH solution is needed to
neutralize 50 ml. of a 0.5N solution of H2SO4?
23. What is the reason for adding lime water to milk that is
to be fed to an infant?
24. NaHCO3 is frequently used to neutralize an acid, (a) Is
it a base? (b) Explain your answer.
25. A glass bottle with a glass stopper contained solid
NaOH. When an attempt was made to open the bottle, it was
found that the glass stopper was firmly cemented to the neck
of the bottle. Explain.
26. In making cream of tomato soup, a pinch of NaHCO3
is added to the tomato puree before the milk is added. Explain
the reason for this practice.
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Take home some pieces of litmus paper and test the fol-
lowing substances for acidic, basic, or neutral properties: vine-
gar, milk of magnesia, ammonia water, table salt solution,
soap solution, sugar solution, and the liquid in an automobile
battery. What do you conclude about the importance of acids
and bases in everyday living?
2. With the help of your chemistry teacher, perform the fol-
lowing experiment: Pour an excess of NH4OH on a solution
of FeCl3. You will get a brown, sticky precipitate of Fe (OH) 3.
Boil thoroughly, then filter. Wash the insoluble Fe (OH) 3
several times with distilled water. What properties of soluble
bases are not possessed by this insoluble base?
3. Using litmus paper as indicator, neutralize (with con-
stant stirring) a known volume of vinegar with some house-
hold ammonia. Report to your class on the relative strength
of these solutions.
SOLUTIONS:
WATER, THE
UNIVERSAL SOLVENT
. . . The sea is the chemist that dis-
solves the mountains and tlie rocks,
pulverizes old continents and builds
new, forever redistributing the solid
matter of the globe. Ralph Waldo
Emerson, 1803-1882
Is chemically pure water found in nature? The purest form of
water in nature is rain water. Although it is perfectly safe for drink-
ing purposes, rain water is not pure water, for it is mixed with the
gases of the atmosphere and with small amounts of dust and other
impurities, which it has washed down from the air.
Eventually all water finds its way to the ocean, the great reservoir.
Rain may fall upon busy streets and flow through sewers to' the sea.
It may fall on the ground and furnish the water so necessary to grow-
ing plants. Water evaporates from the leaf-surfaces of plants and
from the surfaces of streams, lakes, and oceans only to fall again as
rain. Without this water-cycle, life as we know it could not exist.
Water that flows over the ground (surface water} collects fine
particles of solid material. The size and the weight of this material
depend on the speed with which the water flows, for fast-moving wa-
ter can carry much heavier material than a slow-moving stream.
Water that soaks into the ground (ground water] carries almost no
load of this type, since the soil acts as a filter and holds the solid par-
ticles back. But ground water is still not pure water, no matter how
clear and sparkling it may be, for it contains minerals that have been
dissolved out of the soil.
217
218 NEW WORLD OF CHEMISTRY
What is a solution? When a teaspoonful of sugar is dissolved in a
glass of water, the sugar completely disappears from view, and we
say that we have made a solution of sugar and water. The substance
that is dissolved is called the solute; the substance in which the
solute dissolves is called the solvent. Thus, when we make a solution
of sugar and water, the sugar is the solute and the water is the solvent.
If we taste a sample of this solution of sugar and water, we find that
it tastes the same, no matter whether we have taken our sample from
the top or the bottom of the solution. This indicates that a solution
has the same composition throughout; that is, it is homogeneous.
If we take care. not to let any of the water evaporate, the sugar will
not settle to the bottom; the solution will remain the same. Of course,
different amounts of sugar will dissolve in a glass of water — the
solution may be very sweet or it may be only slightly sweet. Thus
a solution differs from a compound, for we can vary the composi-
tion of a solution. The composition of a compound does not vary.
A uniform mixture of solute and solvent that does not conform to
the law of definite proportions is called a solution. The kinetic
theory of matter helps to explain the mechanism of solution. Accord-
ing to this theory, the solute breaks down into molecules which dis-
tribute themselves between the molecules of the solvent. For this rea-
son, a solution is sometimes called a molecular dispersion of a solute
in a solvent. If the solute is colorless (for example, sugar) , it can no
longer be identified by sight when it is in solution.
Difference between dilute and concentrated solutions. A pinch of
salt in a gallon of water makes a very dilute solution. Half a pound
of salt added to the same amount of water makes a concentrated solu-
tion. When only a small amount of solute is dissolved in a solvent, we
have a dilute solution; when the amount of solute dissolved is con-
siderable we have a concentrated solution.
What determines the amount of solute that will dissolve? There
is a limit, of course, to the amount of a solute that will dissolve in
a given volume of a solvent. Several factors determine the amount of
solute that will enter into solution. The most important factors are
the nature of the solvent and the nature of the solute. Large amounts
of certain substances, such as salt, dissolve in water, but the amount
of gold that dissolves in water is extremely minute. Iodine dissolves
only slightly in water, but it is very soluble in alcohol, forming an
alcohol solution known as a tincture. Just why a substance dissolves
in one solvent and not in another is not thoroughly understood.
Temperature has a great deal to do with the amount of a solute
that will enter into solution. More sugar will dissolve in hot tea than
SOLUTIONS: WATER, THE UNIVERSAL SOLVENT
219
in iced tea. In general, most solids dissolve in larger amount in warm
liquids than in cold liquids. This is not true of gases, for the higher
the temperature of a liquid, the smaller the amount of gas that will
dissolve in it. You are familiar with this fact if you have ever noticed
that gas from a warm bottle of soda pop escapes more rapidly than gas
from an ice-cold bottle of the same beverage.
TABLE 9.
SOLUBILITY ^N
CURVES 1 3
Solubility varies with ^
I temperature. A nearly ^ °*
] straight line indicates H §
I that an increase in
I temperature produces
1 a regular increase in
solubility.
0 10 20
30 40 50 60 70
Temperature, °C.
Another important factor that determines the solubility of a gas is
pressure. The higher the pressure, the greater the amount of gas that
will dissolve. For this reason, when we want to make a concentrated
solution of carbon dioxide in water (the carbonated water you see at
every soda fountain) , we add the gas to the water under high pres-
sure and at low temperature. The weight of gas that dissolves is pro-
portional to the pressure on the gas (Henry's law) . In other words,
if we double the pressure, twice as much gas will dissolve.
Two factors that help to determine the amount of solute that will
dissolve in a definite period of time (speed of solution) are the degree
of subdivision of the solute and the extent to which particles of solute
and solvent are closely intermingled by stirring. As you know, the
finer the particles and the more vigorous the stirring, the quicker the
solid will dissolve. However, while these two factors affect the speed
at which a solute dissolves, they do not affect the maximum amount
of solute that will dissolve.
What is a saturated solution? If we add a salt, such as alum, to a
given volume of water at a definite temperature and under fixed
small crystal
of hypo
supersaturated
solution of hypo
(sodium thiosulfate)
saturated
solution
of hypo
3,':
crystals '
starting to form
hypo crystals
Fig. 45. Formation of hypo crystals from a supersaturated solution
7
conditions of pressure, the water will continue to dissolve the salt un-
til it has dissolved a certain amount. After this amount is dissolved,
no more alum will dissolve, and any additional alum that is added
will settle to the bottom of the water and remain there. A solution in
which no more of the solute will dissolve, at that particular tempera-
ture and pressure, is a saturated solution. As long as the solvent will
dissolve more solute, the solution is said to be unsaturated.
The addition of another crystal of the solute to a solution will
indicate saturation or unsaturation. If the crystal added does not
dissolve, the solution is saturated. From the point of view of the
kinetic theory of matter, some of the crystals added to a saturated
solution do dissolve, but just as many molecules of the solute come
out of solution. Thus the crystals added do not appear to dissolve.
How supersaturation is used in purifying solids. It may sound
contradictory to say that it is possible to prepare a supersaturated
solution, that is, a solution which contains more of the solute than a
saturated solution. But it is possible. For example, to prepare a super-
saturated solution of hypo (Na2S,O8 - 5H,O) , we first make a satu-
rated solution of this salt in boiling water, and then slowly cool the
solution. The excess hypo does not come out of solution, as we would
expect, but remains in solution. Because this solution contains more
solute than it normally holds at the lower temperature when satu-
rated, it is called a supersaturated solution.
Supersaturation is an unstable condition, and if the solution is
disturbed by adding a tiny crystal of the solute, all of the excess salt
separates out and the solution becomes saturated. Since only pure
020
SOLUTIONS: WATER, THE UNIVERSAL SOLVENT 221
crystals will separate out, this method is often used commercially in
preparing chemically pure (C.P.) salt crystals. Because impurities
are present to a small degree as a rule, they separate out later.
The effect of a solute on freezing and boiling points. As you know,
the addition of salt to water lowers the freezing point of the solu-
tion. In general, a solute raises the boiling point and lowers the
freezing point of a solvent a definite number of degrees. This prin-
ciple is used in preparing antifreeze mixtures, which freeze at tem-
peratures much lower than 0°C., the freezing point of pure water.
A water solution that is 50 percent grain alcohol does not freeze un-
til it reaches — 31°C. Hence, solutions containing alcohol or other
solutes are used in automobile radiators in cold weather.
All substances do not have the same effect on freezing and boiling
points. Certain acids, bases, and salts raise the boiling point and
lower the freezing point two or three times as much as other sub-
stances including sugar, alcohol, and glycerin. This singular behavior
has an important relation to dissociation (see Chapter 16) .
The difference between a solution and a suspension. If a substance
does not actually dissolve but becomes mixed somewhat uniformly
with a solvent and then separates on standing, the mixture is not
a true solution but a suspension. Fine clay or other fine materials
intimately mixed with water is a suspension.
The chief differences in behavior between a true solution and a
suspension depend upon differences in the size of the particles of the
solute in solution compared with the size of the particles in suspen-
sion. Particles of solutes in solution are in the molecular state of
subdivision. But even the finest particles in suspension are much
larger than molecules; hence, they settle out.
Particles in suspension may be separated by filtration; those in a
solute cannot be separated by filtration. The solute of a true solution
has a greater effect upon the boiling and freezing points of the
solvent than the particles in suspension have on a suspension.
In the borderland between true solutions and suspensions is an-
other condition of matter called the colloidal state. This condition
of matter is so important that all of Chapter 38 is devoted to it.
How we purify water. Even before the germ theory of disease
was proved, several methods were in use for the purification of
water. In ancient times, water was made more fit to drink by filter-
ing, boiling, or by allowing suspended impurities to settle out. Laws
were passed to guard against the contamination of river water used
for drinking purposes by prohibiting the washing of clothes and the
disposal of refuse and sewage in it.
222 NEW WORLD OF CHEMISTRY
In a chemical sense most drinking water is not pure, tor it con-
tains dissolved minerals and gases. It is pure in a biological sense,
which means that it is comparatively tree from bacteria and other
organic matter harmful to health.
Water used for drinking purposes may be put through one or
more of the following processes of purification, depending upon the
nature of the impurities it contains:
(1) Aeration consists of spraying the water into the air or letting
it tlow down a series of steps in thin sheets so that sunlight and oxy-
gen may kill most bacteria present. (2) Filtration consists of strain-
ing the water through a suitable sieve (filter) , thus separating sub-
stances either in suspension or afloat. Sand filtration dates as far back
as 1829, when London purified its drinking water by passing the
water through beds of line sand. (3) Chlorination, discussed in
Chapter 10. (4) Ozonalion, discussed in Chapter 2. (5) Coagulation
consists of adding alum or other chemicals that cause the organic
matter containing bacteria and other suspended material to settle
out. Coagulation is discussed more fully on page 396. (6) Chlor-
amination, another process of water purification using both chlo-
rine and ammonia is discussed on page 253. (7) Demineralization
by means of synthetic resins (see page 191) .
Modern sewage disposal. Closely related to the problem of main-
taining an adequate supply of safe drinking water is the problem of
disposing ol the domestic and industrial wastes that are produced
especially under the crowded living conditions of modern cities.
Sewage, as these wastes are called, includes chiefly the organic
wastes that may be disposed of by ft nl re faction, a process that con-
sists of a combination of bacterial action and oxidation. In some
communities, sewage is discharged directly into streams which carry
the sewage away. The dissolved oxygen in the water of the stream
eventually oxidizes sewage; but, as a result, the dissolved oxygen
Fig. 46. Water filtration. As the water passes
through the filtering layers, solid matter and
many germs are removed.
water
film
sand
fine
tile
cement
pipe
An effective and relatively
inexpensive method of puri-
fying water is to aerate it
by spraying it info the air.
is used up and all the fish and other higher plant and animal life
normally found in the stream are unable to live in the water because
they cannot obtain oxygen. Only a few of the lower forms of life,
such as certain algae and bacteria, can live in the polluted stream,
and the water is unfit for almost all purposes. Oxygen-consuming
factory wastes are sometimes handled in disposal wells. Fortunately,
the oxygen of the air cleanses the polluted stream in the course of
its meanderings. The distance necessary to cleanse the stream de-
pends upon several factors — the amount of sewage discharged, the
size of the stream, its rate of flow, and so forth.
Most modern cities dispose of sewage by more modern and less
harmful methods of treatment such as the activated sludge process.
The sewage, which flows to the modern treatment plant in great
quantities of flushing water, is first run into closed tanks where the
solids settle as a sludge, or is run through sieves and screens that
remove the suspended solids. Certain bacteria present in the sew-
age decompose the sludge, liberating both nitrogen and methane
gas, and convert the sewage into a nontoxic, humus-like waste that
may be used as a fertilizer. The partially purified water may be
further purified by one of the methods of purification already dis-
cussed, or it may be discharged into a neighboring stream or body of
water. In certain types of installations, the methane gas produced
during the digestion of the sludge is burned in gas engines to pro-
duce the energy necessary to operate the sewage-treatment plant.
How water can be made chemically pure by distillation, None
of the methods mentioned removes all impurities from water. Not
one of them completely removes substances dissolved in water. Dis-
solved substances are completely removed from solution both in
industry and in the laboratory by distillation. Because the boiling
points of substances differ, it is possible to separate solute from
223
thermometer
inlet from
faucet
--distillate
Fig. 47. Distillation of water using a Liebig condenser. Why should
cold water enter the jacket at the lower end?
solvent. In general, solids that are dissolved in water have higher boil-
ing points than water, and remain behind after the water has been
boiled off. Certain liquids, including glycerin, have higher boiling
points than water; others, as alcohol, have lower boiling points.
Hence two or more liquids in solution may be separated by distilla-
tion. Gases in solution are driven off soon after the water is heated.
During distillation, water is first boiled, or evaporated, and
then the steam, or water vapor, is cooled. This cooling condenses it
into water again. Distillation is thus a double process, including both
evaporation and condensation. The first portion of the distillate, the
liquid that results from the condensation of the vapor, may contain
small amounts of dissolved gases. The final portion of the distillate
may contain small amounts of liquids or even dissolved solids whose
boiling points are close to that of water. If these two portions are dis-
carded, the rest of the distillate will be free of all impurities.
Stills used for industrial purposes are made of such material as
copper, steel, lead-lined steel, and fused silica (SiO2) .
How water is distilled in the laboratory. The laboratory appara-
tus for water distillation consists of a flask, a Liebig condenser
made popular in 1850 by Justus von Liebig (ton le'biK) , a German
chemist famed for his contributions to organic and agricultural
chemistry, and a receiver. Impure water, which, for the purpose of
the experiment, may contain small amounts of ammonia, salt,
and red ink, is boiled in a sidearm flask, as shown in Fig. 47. A
thermometer indicates the boiling point of the solution. The steam
and water vapor pass into the inner glass tube of the Liebig con-
denser, which is surrounded by a glass jacket having a glass inlet
and outlet for water. To condense the water vapor before it escapes
from the inner glass tube, cold water is circulated through the outer
tube of the condenser. As a result of this cooling, distilled water
collects as the distillate.
Fractional, partial vacuum, and high vacuum distillation. When
the boiling points of the impurities are very close to the boiling
224
SOLUTIONS: WATER, THE UNIVERSAL SOLVENT
225
point of the solvent, simple distillation is not very effective. For ex-
ample, the boiling point of grain alcohol is 78°C., and if separated
from water by a single distillation, the distillate is not pure.
Oil refiners face a difficult problem in separating the liquids that
make up crude petroleum. The boiling points of these liquids are
so close together that he must resort to fractional distillation. This
consists of heating the mixture carefully and separating it into its
various liquid fractions, which boil off at different temperatures. For
example, the first fraction might boil off below 200°C., the next be-
tween 200°C. and 220°C., and so forth. In each case, it is necessary
to purify each fraction further by additional distillations.
It often happens that liquids cannot be separated by fractional
distillation, because the high temperature required may char or
otherwise decompose the solute. In such cases partial vacuum, or
even high vacuum distillation is used. By reducing the pressure over
the solution, its boiling point is also lowered.
According to the kinetic theory of matter, the molecules of water
leave the liquid and become water-vapor molecules much more
easily when the pressure over them is decreased. Water, for exam-
ple, under normal atmospheric pressure boils at 100°C. If this pres-
sure is lowered to one-half normal, water will boil at 82°C., and
the process of distillation can be carried through at this lower tem-
perature. Thus, water is removed from milk in making evaporated
milk by partial vacuum distillation. Water is removed from sugar-
cane sap in making granulated sugar by reduced pressure distilla-
tion. High vacuum distillation is used in making vaccines, serums,
antibiotics, blood plasma, frozen orange juice, and for coating metals.
The principle of distillation is used in both the huge petroleum
"cracker" (left) and the laboratory fractionating equipment (right).
Shell Oil Company
226 NEW WORLD OF CHEMISTRY
Distilled water has many uses. Because distilled water is free from
dissolved gases from the air, it is flat and insipid in taste. Its prop-
erties are those of chemically pure water (see page 65) . Ordinarily,
distilled water is not used for drinking purposes on land. However,
at sea, drinking water is commonly prepared from sea water by dis-
tillation. The economic production of drinking water from sea wa-
ter, which contains about 3.5 percent of dissolved salts, is very old,
dating back at least to the time of Aristotle (about 350 B.C.) .
Today, ships of nearly all kinds obtain drinking water by distill-
ing sea water under reduced pressure. Seagoing vessels also produce
water for use in boilers by distilling sea water under reduced pres-
sure. In this way, all salts are removed, and subsequent treatment re-
moves all traces of dissolved gases. Consequently, the operation of
the boilers is not reduced in efficiency by the formation of large
quantities of boiler scale (see page 489) .
Today, aviators who are forced down at sea, and all abandon-ship
equipment of warships and merchant vessels, carry small kits that
may be used to prepare safe drinking water from sea water easily
and simply. Such kits produce drinking water by chemical means.
Their chemistry is discussed on page 491.
Distilled water is used in storage batteries, in which ordinary drink-
ing water should not be used because of the harmful effect of its im-
purities on the plates of the battery. Distilled water is indispensable
in many scientific and industrial operations. In the preparation of
vaccines, pharmaceuticals, certain dyed textiles, and C.P. chemicals,
distilled water is commonly used.
Heavy water, a powerful tool of research. The heavier hydrogen
isotopes, deuterium and tritium (see page 56) , may be represented
respectively by the symbols D and T. It has been determined that the
nucleus of deuterium contains one proton and one neutron. There-
fore its atomic weight is two. The nucleus of tritium contains one
proton and two neutrons and it has an atomic weight of three.
When water is formed from oxygen and one of the two heavy hy-
drogen isotopes, its molecular weight is greater than 18, the molecu-
lar weight of ordinary water. This water is called heavy water. Heavy
water can also be formed from deuterium or tritium and any one
of the three heavy isotopes of oxygen. Deuterium oxide, D2O, is
present in ordinary water to the extent of about one part of D2O
in 5000 parts of water. Tritium is seldom found in nature but is or-
dinarily made in an atomic pile.
Deuterium oxide has been prepared in large quantities. It differs
to a small extent from ordinary water in both freezing and boiling
SOLUTIONS: WATER, THE UNIVERSAL SOLVENT 227
points. Its maximum density occurs at 11.6°C. rather than at 4°C.,
as is the case with H2O. Heavy water is used in atomic piles as a
moderator.
Deuterium oxide appears to arrest the growth of seedlings; tad-
poles die prematurely in it; it is, however, not toxic to man and the
higher animals.
Heavy hydrogen has been substituted for ordinary hydrogen in
certain fats, and the course and changes which these "tagged" fat
molecules have undergone on their way through the animal body
have been studied by the tracer technique.
YOU WILL ENJOY READING
Cerna, Wendell W. "Industrial Water Conditioning Proc-
esses/' Journal of Chemical Education, March, 1943, pp. 107-
115, and April, 1943, pp. 191-197.
Ellis, Cecil B. Fresh Water from the Ocean. Ronald Press
Company, New York, 1954. This is a conservation study deal-
ing with water for cities, industry, and irrigation.
Goldblatt, L. A., Ed. Collateral Readings in Inorganic Chem-
istry. D. Appleton-Century Co., New York, 1937 (2nd series,
1942). No. 8 of the 31 articles in this collection deals with
"Factors Contributing to Quality of Public Water Supplies."
Written by H. E. Jordan.
USEFUL IDEAS DEVELOPED
1. A solution is a uniform mixture of solvent and solute
which does not conform to the law of definite proportions.
2. A dilute solution contains very little solute in comparison
with the solvent; a concentrated solution contains a large
amount of solute.
3. Some of the factors which determine the amount of solute
that will dissolve in a solvent are: (1) the nature of solute
and solvent, (2) temperature, and (3) pressure. The speed of
solution depends upon (1) the state of subdivision of the
solute, and (2) how thoroughly the solute and solvent are
intermingled.
4. All solutes raise the boiling point and lower the freezing
point of the solvent. The amount of solute determines the
number of degrees of change. Acids, bases, and salts in solution
affect boiling and freezing points to a greater degree than do
other substances, including alcohol, sugar, and glycerin.
228 NEW WORLD OF CHEMISTRY
5. Solutions and suspensions differ in the following respects:
(1) In a suspension, the mixture separates on long standing.
(2) The particles in suspension are much larger than the
particles of a solute in a solvent (which are in the molecular
state of subdivision) . (3) The particles of a suspension may
be separated out by filtration; the particles of a solution can-
not be so separated. (4) A solute has greater effect on the
freezing and boiling points than has a material in suspension.
6. The colloidal state of matter is in the borderland be-
tween true solutions and suspensions.
7. Distillation consists of evaporation and condensation.
Impurities can be removed by distillation because of the dif-
ference in the boiling points of a solvent and its solute.
USING WHAT YOU HAVE LEARNED
Group A
1. (a) What is the purest form of H2O found in nature?
(b) What property of H2O makes it almost impossible to find
pure H2O in nature?
2. (a) What is a solution? (b) In a solution of NaCl in
H2O, which is the solvent and which the solute? (c) What is
a tincture?
3. Distinguish a dilute solution from a concentrated so-
lution.
4. Name four factors that determine the degree and the
speed of solubility of a substance.
5. Why is more sugar or more stirring required to sweeten
iced coffee than hot coffee?
6. Why should sealed bottles of H2O heavily charged with
CO2 be kept cold?
nr
7. How would you proceed to prepare a saturated water
solution of washing soda, Na2CO3?
8. How could you change a saturated solution of CuSO4 to
an unsaturated solution having the same volume?
9. Without changing the temperature or volume of a solu-
tion of CO2 in H2O, how could you change an unsaturated
solution of this gas into a saturated one?
10. How could you determine whether a solution is satu-
rated, unsaturated, or supersaturated?
11. Mention one commercial use of supersaturation.
SOLUTIONS: WATER, THE UNIVERSAL SOLVENT 229
12. What is the effect of a solute on the freezing and boiling
points of a solvent?
13. What principle is involved in the use of an antifreeze
mixture in an automobile radiator?
14. How do true solutions and suspensions differ?
15. What is the name given to the state of matter that is on
the borderland between true solutions and suspensions?
16. What are seven methods used to purify drinking water?
17. How is sewage disposed of in modern sewage-disposal
plants?
18. What may be the results of ineffective methods of sewage
disposal?
. . I . . .
I
19. Which method or methods of purifying H2O produce
chemically pure H2O?
20. (a) What is distillation? (b) Upon what fact does the
separation of impurities by distillation depend?
21. Why is the middle portion of a distillate the purest?
22. Make a diagram of the apparatus used in the laboratory
distillation of H2O.
23. (a) Is the drinking water of a large city, such as Chicago
or New York, distilled? (b) Give reasons for your answer.
24. Petroleum is refined by fractional distillation. Why?
25. Why are liquids often distilled under reduced pressure?
26. State the physical properties of distilled water.
27. For what is distilled water used?
28. What are the desirable characteristics of drinking water?
29. What is the difference between wholesome water and
chemically pure water!
30. (a) What is heavy water? (b) How does it differ from
ordinary water? (c) Name two uses of heavy water.
Group B
31. Devise an experiment to show that perfectly clear spring
water contains impurities.
32. Small filters attached to household faucets sometimes
become a menace to health. Explain this statement.
230 NEW WORLD OF CHEMISTRY
33. What would be the effect of the continued use of rain
water in storage batteries?
34. Is water obtained by melting ice from a lake purer than
the water of that lake? Explain your answer.
35. If a liquid is colorless, odorless, and clear, how could
you determine whether it is a solution or a pure compound?
36. How does a solution of NaCl in water differ from a mix-
ture of NaCl and sugar?
37. Can a dilute solution be a saturated solution also? Ex-
plain.
38. Explain the operation and principle of the pressure
cooker.
39. (a) Explain solution by means of the kinetic theory of
matter, (b) According to the kinetic theory, explain why it is
easier to evaporate or distill a liquid under reduced pressure.
40. Would there be any advantage or disadvantage in
making solutions of (a) table salt, (b) calcium hydroxide,
with hot water instead of cold? (c) Explain each answer.
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Visit your local sewage-disposal plant, and study its op-
eration. Make a report on this visit to your class.
2. Obtain some Rochelle salt (U.S.P.) from your teacher
or druggist and prepare pure crystals from it. Bring them to
class and explain the process you used. (As a substitute, you
may prepare "rock candy" crystals.)
3. Harold C. Urey, discoverer of heavy water, in a lecture in
1938, made the following statement, "I believe I speak for the
vast majority of all scientific men. Our object is not to make
jobs and dividends. These are a means to an end, merely inci-
dental. We wish to abolish drudgery, discomfort and want
from the lives of men, and bring them pleasure, comiort,
leisure and beauty. Often we are thwarted but in the end we
will succeed." Write a short report either for or against this
view, or organize a class discussion on this subject.
4. Study the methods of water purification in your own
community. Report to your class on this subject. Find out, if
you can, the mineral content of the water and the amount and
kind of suspended solids present both before and after treat-
ment.
5. Using hypo (sodium thiosulfate) prepare 10 ml. of a
supersaturated solution. Allow the solution to cool slowly to
room temperature. Then add a crystal of hypo. Repeat this
before your class if it can be arranged.
IONS
AND DISSOCIATION
. . . / heard Cleve say: "Do you
believe sodium chloride is dissolved
into sodium and chlorine? In this
glass I have a solution of sodium
chloride. Do you believe there are
sodium and chlorine in it?" "Oh,
yes," Ostwald replied, "there is some
truth in that idea." . . . Cleve threw
a look at Ostwald which clearly
showed that he did not think much
of his knowledge of chemistry.
Svante Arrhenius, 1925
Two eternal questions: How? and Why? Science is constantly try-
ing to answer two questions — how and why. Often it is not too dif-
ficult to answer the hows but the whys — well, that is a different
story. Theories must be formulated, tested, and adapted to keep
them in accord with all the observed facts, and that is a big job.
At least two great theories underlie much of chemistry: the elec-
tron theory and the theory of dissociation. In a sense, the theory of
dissociation is but an aspect of the electron theory, for the theory
of dissociation is explained in terms of the electron theory. These
two great theories explain some of the hows and whys of chemistry.
A Swedish boy tackles some puzzling questions. Why is distilled
water a nonconductor of electricity? Why do water solutions of some
substances conduct electricity yet water solutions of other substances
do not? What causes some acids and bases to be strong while others
are weak?
These were some of the unsolved problems that confronted
chemists when Svante Arrhenius (ar-ra'ni-us) was still at school in
Sweden.
Arrhenius not only wondered about these problems but set to work
to solve them. He had some unusual notions of his own about the
231
232 NEW WORLD OF CHEMISTRY
way in which electricity passes through solutions. Day after day, and
often far into the night, he worked in his laboratory with hundreds
of different solutions. For two years he toiled ceaselessly.
Arrhenius attempted to formulate a theory that would explain
what he had observed. In those days his whole world, both of wak-
ing and sleeping hours, was one of solutions, electric currents, and
formulas. For him the rest of the world did not exist. One night he
sat up very late. Suddenly, like a flash, he saw the answer to the great
riddle. "I got the idea/' he wrote, "in the night of the 17th of May
in the year 1883, and I could not sleep that night until I had worked
through the whole problem."
He went to his teacher of chemistry. "I have a new theory of
electrical conductivity/' Arrhenius told him. The professor looked
at this boy and said: "You have a new theory? That is very interest-
ing. Good-by." But Arrhenius did not lose heart. He wrote to the
leaders in chemistry. Most of them were hostile to his revolutionary
theory. After a long struggle, however, they were forced to admit its
probable truth and later saw Arrhenius awarded the Nobel prize in
chemistry. And so it happened that a mere boy, with the clear insight
and the creative imagination of a truly great scientist, stepped in and
cleared away an obstacle that had stood squarely in the path of chem-
ical progress.
How Arrhenius explained the conductivity of solutions. The the-
ory proposed by Arrhenius is known as the theory of ionization, or
dissociation. Arrhenius assumed that when an electrolyte, such as
sodium chloride, dissolves in water, it tends to dissociate, or ionize.
That is, it tends to break apart into electrically charged atoms or
groups of atoms (radicals) . Arrhenius used the term ion, meaning
wanderer, to refer to an atom or group of atoms carrying an electric
charge. He represented the dissociation of the sodium chloride mole-
cule into ions when it dissolves in water thus:
NaCl — » Na+ (sodium ion) + Cl~ (chloride ion)
He represented the dissociation of the sodium nitrate molecule thus:
NaN08->Na++(NO3)-
Arrhenius could not see the ions in solution. They are far too small
to be seen. He advanced the theory that they were present because
he could account for what he observed only by assuming such ions to
be present. To him, when an electrolyte dissolves, a certain number
of its molecules immediately split up into ions. Thus an electrically
IONS AND DISSOCIATION 233
neutral compound tends to dissociate into a number of positively and
negatively charged particles, or ions. They move about in all direc-
tions until the passage of an electric current draws each ion to the
electrode bearing an opposite charge.
It is the ions that carry an electric current, or flow of electrons.
Hence a substance that does not dissociate into ions is a nonelectro-
lyte, and a substance of which a large proportion dissociates in water
solution is a good electrolyte. Arrhenius' theory of ions and of elec-
trolytic dissociation is today, with certain modifications, universally
accepted as the correct explanation of the conductivity of solutions.
How an ion differs from an atom. There are two main differences
between an atom and an ion: (1) An atom is electrically neutral; an
ion is positively or negatively charged. (2) An atom always consists
of a single element; an ion may consist of one or more than one ele-
ment, as in the case of the ammonium ion, (NH4) +. A sodium atom
is quite different from a sodium ion. The former is a silvery metal
particle that reacts violently with water; the latter is a colorless par-
ticle that has no noticeable reaction with water.
The dissociation of zinc chloride may be represented thus:
Note the change from C12 in ZnCl2 to 2 Cl~ when the zinc chloride
is dissociated into ions. The number of positive or negative charges
on the ion is equal to the valence of the element or radical.
PRACTICE WORK ON DISSOCIATION EQUATIONS. Write equations
showing what happens when the following electrolytes dis-
sociate:
a) Nitric acid, HNOr g) Sulfuric acid, H2SO4.
b) Lithium hydroxide, LiOH. h) Potassium carbonate,
c) Potassium hydroxide, K2CO3.
KOH. i) Sodium phosphate,
d) Ammonium hydroxide, Na3PO4.
NH4OH. j) Potassium chlorate, KC1O3.
e) Barium hydroxide, k) Magnesium bicarbonate,
Ba(OH)2. Mg(HC08)2.
f) Aluminum chloride, A1C13. 1) Calcium nitrate,
Ca(N03)2.
Where and how are free ions formed? Most commonly, dissocia-
tion takes place in a water solution. However, it also occurs in com-
pounds such as sodium hydroxide, NaOH, when they are heated un-
til they melt, or fuse. Such compounds are made up, not of molecules,
as was previously supposed, but of ions in a definite pattern (see
234
direction of flow
of electrons
heated filam
cathode
Fig. 48. Simplified diagram of a radio
tube showing the hot filament emitting
a stream of electrons.
page 631). As the compound is heated, its ions vibrate more rap-
idly and eventually tear apart, leaving them free to carry an electric
current. Gases, too, at very high temperatures dissociate into ions,
thus becoming conductors of electricity. Flames, x-rays, and radio-
active elements ionize the air around them.
A metal, when heated to incandescence, boils off electrons from
the outer rings of its atoms. This principle is used in the electron
tubes so familiar to us in radios, television sets, and other electronic
instruments. Certain types of electron tubes are filled with a gas or
mixture of gases at low pressure. Electrons from the hot cathode of
the tube strike the molecules of the gas, ionizing them.
The positive ions move toward the negative electrode, or cath-
ode, and the negative ions and electrons move toward the positive
electrode, or anode. In a radio tube, the anode is the plate, and
the cathode may be the hot filament. Thus a stream of electrons or
electric current flows through the tube. Because a gas-filled tube
contains more molecules that can be ionized than does a vacuum
tube from which most of the molecules have been removed, relatively
large currents can flow through a gas-filled tube. The thyratron tube,
often employed as a rectifier, is a common example of a gas-filled
electron tube.
How Arrhenius explained the action of acids. All acids in water
solution contain free hydrogen ions. These hydrogen ions determine
the typical properties of acids. It is the free hydrogen ion that turns
blue litmus pink, has a sour taste, is replaceable by a metal, and
neutralizes a base by combining with its hydroxyl ion.
The number of hydrogen ions present determines the strength of
an acid. A strong acid is one that dissociates easily, and thus pro-
duces a large number of hydrogen ions. The most commonly used
strong acids are hydrochloric, nitric, and sulfuric acids. Carbonic
acid, H2CO3, sulfurous acid, H2SO3, and boric acid, H3BO8, are weak
acids, because they do not dissociate easily, and thus form only a
small number of hydrogen ions in solution.
IONS AND DISSOCIATION 235
How Arrhenius explained the action of bases. What has been
said about strong and weak acids refers equally well to strong and
weak bases, except that the determining factor here is the number of
free hydroxyl (OH) ~ ions present. In water solution, a strong base,
such as sodium hydroxide, NaOH, potassium hydroxide, KOH, or
lithium hydroxide, LiOH, forms large numbers of hydroxyl ions.
A weak base, such as ammonium hydroxide, NH4OH, produces a
comparatively small number of (OH)- ions.
Some bases, like copper hydroxide, Cu (OH) ,, and aluminum hy-
droxide, Al (OH) 3, are extremely weak, because, in addition to the
fact that they form a relatively small number of ions, they are also
only sparingly soluble. The properties of bases are caused by the pres-
ence of free hydroxyl ions. Hence dry, solid sodium hydroxide has no
basic properties; that is, it does not react as a base.
How the theory of Arrhenius explained neutralization. The hy-
drogen ion and the hydroxyl ion are responsible for acidic and basic
properties respectively. If these free ions are removed, acidic and
basic properties are destroyed. This is exactly what happens in neu-
tralization. The hydrogen ion of the acid unites with the hydroxyl
ion of the base, forming water, ;i nonelectrolyte.
Na+ + (OH)~ + H+ + Cl~ -> Na+ + Cl~ + HOH
Hydrogen-ion concentration (pH) and how it is measured. It is
frequently important to know whether complete neutralization has
been produced, or what the degree of acidity or alkalinity (basic-
ity) of a substance is. Sugar-refiners, brewers, paper-makers, electro-
platers, sanitary engineers, and bacteriologists must have a working
knowledge of the acidity or alkalinity of many substances.
Chemists make use of a hydrogen-ion scale in which the unit is
the /;H value (pronounced by reading the letters p, H) , just as the
unit of temperature is a degree. A jm value of seven is considered
Chemists in professional labora-
tories use special equipment to
determine the pH value of sol-
utions.
236
NEW WORLD OF CHEMISTRY
true neutralization. Pure water has such a value. A pn value less than
seven indicates an acid condition. Thus, saliva has a pn value of
about 6.9 and is slightly acid. A pn value greater than seven indi-
cates an alkaline condition. Thus, normal blood has a pu value of
about 7.3 and is slightly alkaline. A value higher than 7.5 repre-
sents a condition of alkalosis. Since the human blood is normally
slightly alkaline, a pn below 7.2 indicates a condition of acidosis.
A pn value of five represents an acidity 10 times as great as a pn
of six.
Litmus may be used as an indicator only with those solutions
whose pn is not less than four or greater than eight. Other indi-
cators have wider or narrower ranges. By comparison with standard
color tubes for each indicator, the degree of acidity or alkalinity
of a liquid can be determined accurately.
Several types of indicator papers have been developed for deter-
mining the pn values of solutions. The use of such papers contain-
ing water-soluble dyes is very convenient and is much quicker than
referring to standard color tubes. One such paper is known as
"Hydrion Paper."
How the degree of dissociation was determined. We have seen
that some substances produce a large percentage of ions and other
substances produce a small percentage of ions. In other words, com-
pounds differ in the extent to which they form ions in solution, that
is, in their degree of dissociation. The degree of dissociation of a
solution depends upon (7) the solute, (2) the solvent, (3) the con-
centration of the solution, and (4) the temperature.
The degree of dissociation is measured by electric conductivity,
that is, the ease with which an electric current passes through a
solution. This may be determined roughly by setting up an appa-
ratus as shown in the illustration below (Fig. 49) . When the bottle
contains carbonic acid (a weak acid) the electric-light bulb glows
faintly. When it contains a dilute solution of hydrochloric acid (a
strong acid) , the light bulb glows brightly, showing that little resist-
ance is being offered to the passage of the current. A solution of
sugar in water, when placed in the bottle, does not produce a glow
in the bulb, and thus shows that a sugar solution is a nonelectrolyte.
Fig. 49. Laboratory setup for studying
the conductivity of a liquid. When will
the bulb light?
platinum electrodes
liquid to be tested
IONS AND DISSOCIATION 237
Arrhenius determined the apparent percentage of dissociation of
many compounds by this method. He found that in one-tenth nor-
mal water solution (0.1N) the apparent dissociation of hydrochloric
acid, nitric acid, and sodium hydroxide was about 92 percent (strong
electrolytes) ; that of sulfuric acid about 61 percent; that of potas-
sium chloride about 86 percent; and that of acetic acid, ammonium
hydroxide, and mercuric chloride only about 1 percent (weak elec-
trolytes) . Later it was found that this method had certain limitations.
How Arrhenius explained abnormal boiling and freezing points
of solutions. The theory of ionization also explains why certain
ac ids, bases, and salts raise the boiling points of their water solu-
tions to an abnormally high degree, though sugar does not to the
same extent. The acid dissociates and produces two or three times
as many particles (ions) as there are molecules of undissociated sugar.
The higher percentage of dissociation of the acid produces a greater
increase in the boiling point, since it is the actual number of par-
ticles (ions or molecules) in solution that determines both the boil-
ing and freezing points of a solution.
How the theory of ionization fits in with the electron theory.
When Arrhenius proposed his theory of dissociation, the electron was
still undiscovered, and the electron theory of matter had not yet
been formulated. Ions, however, fit in beautifully with our present
electron concept of matter. Atoms become ions by gaining or losing
electrons. For example, if an electrically neutral atom of sodium
(Na±) loses 1 electron, it becomes a positively charged particle, and
we represent it Na+. This is a sodium ion (Na° — € — > Na+) . Sim-
ilarly, if an electrically neutral atom of chlorine (Cldb) gains I elec-
tron, it becomes a negatively charged particle, and we write it Cl~.
This is a chloride ion (Cl° + € -> Cl~) .
The compound made up of these elements (NaCl) is really com-
posed of both of these ions held strongly together, making what is
known as an ion-molecule that is electrically neutral. However, when
these ion-molecules are placed in water, they are split into two parts,
and so long as water is present, we have separate sodium ions and
chloride ions. Dissociation, then, may be represented as shown below.
Fig. 50. Dissociation of a sodium chloride molecule, ac-
cording to the electron thepry.
NaC1 -molecule *- Na+ (ion) + CT (ion)
(+11-11-0) (-H7-17-0) (+11-10-+1) (+17-18--1)
238 NEW WORLD OF CHEMISTRY
The charge on the ion is thus seen to be the same as the valence of
the element.
How the theory of Arrhenius was later modified. In general,
Arrhenius' theory has stood the test of time very well. In three re-
spects, however, it has been modified slightly.
1) The Swedish chemist thought that water simply kept the ions
apart, but today we have a better understanding of the impor-
tant role of the solvent. The covalent water molecule is pictured
as a tiny magnet with a plus hydrogen end and a less positive oxy-
gen end. Such a molecule is called a dipole.
When crystals of the polar compound, NaCl (made up of a
lattice of Na and Cl ions — see page 630) , are added to water, the
positive end of the dipole molecule attracts the negative Cl ion;
the other end of the dipole attracts the positive Na ion. The Na
and Cl are thus dissociated and free sodium and chloride ions are
produced in solution. When the water is evaporated, the ions re-
combine forming ion-molecules of NaCl crystals. NaCl and other
ionic compounds are strong electrolytes.
2) The Br0nsted-Lowry theory has added to what Arrhenius be-
lieved about the ionization of acids. According to this theory, an
acid is a proton (hydrogen ion) donor; that is, any compound
that tends to lose a proton (H+) to another substance is an acid.
For example, in the case of a water solution of HC1, the water
molecule combines with the positive hydrogen ion end (proton)
of the HC1, forming a hydrogen (hydronium) ion, leaving the
chlorine negatively charged, thus:
H2 O + H Cl ^H20 . H+ + Cl"
V^^CI) hydronium ion
As HC1 dissolves in water, heat is liberated, indicating a chemical
change is taking place. The hydronium ion is an acid since it may
give up its proton. It is also written as H8O+.
A base is a proton acceptor. A base is any substance which com-
bines with a hydrogen ion or proton. Thus the water molecule
which will accept a proton is a base, though not an active one.
Bases are often negative ions such as OH~, which is an active one.
3) Insofar as weak electrolytes are concerned, the theory ex-
pressed by Arrhenius is still correct. They are not completely dis-
sociated. In the case of strong electrolytes, however, it is now be-
lieved that they dissociate completely. The fact that the heat of
4(OH)-*-2H2
hydrogen^ ^JB^Wiiii^^Xoxygen
Fig. 51. Action during the
electrolysis of water.
neutralization of all strong acids and bases is the same is one evi-
dence in support of this belief. The fact that when dilute solu-
tions of strong electrolytes are mixed there is no trace of heat is
further proof. The degree of dissociation of weak electrolytes can
be determined accurately, but that of strong electrolytes cannot
because of the disturbing electrical effects of a large concentration
of ions. In such a high concentration, ions do not behave as "free
ions" and only an apparent ionization is obtained.
The theory of dissociation as stated by Arrhenius is still useful.
Other new ideas have added to the accuracy of our concept of dis-
sociation. For example, we know that molecules of compounds such
as carbon dioxide do not dissociate in solution and are nonelectro-
lytes because they are nonpolar. However, these new ideas are of
value chiefly in dealing with phenomena which are beyond the scope
of this book. For this reason, we shall follow the theory of ioniza-
tion as Arrhenius originally presented it.
What happens during the electrolysis of water? So far we have
considered pure water a nonelectrolyte. Roughly speaking, this is
true, since in about 1 billion grams of water only about 1 gram of
free hydrogen ions is present. Yet this very slight dissociation of wa-
ter [H,O — » H+ + (OH) -] is important. Water is such a poor con-
ductor that acid must be added to conduct the current. Nevertheless,
the hydrogen and oxygen that are products of electrolysis come from
the water. How?
The positive hydrogen ions travel to the negative cathode. The
cathode, connected to a battery, is supplied with electrons. When the
hydrogen ion reaches the cathode, an electron, e, from the cathode
is given up to the hydrogen ion, which changes to a neutral hydro-
gen atom. This atom immediately joins another hydrogen atom and
is liberated as a molecule of gas, thus:
H+ + c -» H°; H° + H° -» H2 T
239
240 NEW WORLD OF CHEMISTRY
The negative sulfate ions, and hydroxyl ions from the water, travel
to the positive anode. The hydroxyl ion gives up its electron more
readily than does the sulfate ion. Having lost its electron, it becomes
the hydroxyl radical, which breaks down, forming water and oxygen.
The sulfate ions remain in solution. The concentration of the acid
remains unchanged.
(OH)--€-»(OH); 4(OH)->2H20 + 02!
What is hydrolysis? We should expect neutral salts to show
neither acid nor basic properties, since they contain neither hydrogen
nor the hydroxyl radical which might form ions. Yet when we add
blue litmus to a solution of copper sulfate, the litmus turns pink, in-
dicating an acid. Where are the free hydrogen ions to account for
this behavior? The answer lies in the fact that water is slightly dis-
sociated.
Cu++ + (SO4)~ + 2HOH «=» 2H+ + (SO4)~ + Cu(OH)2
Since the copper hydroxide formed is only very slightly dissoci-
ated, there will be some union of Cu++ and (OH) -, thus forming this
very weak base, and liberating an excess of hydrogen ions that form
the strong acid, sulfuric acid. Therefore, the solution* is slightly acid.
This reaction is the reverse of the neutralization of sulfuric acid by
the base Cu (OH) ,. Hydrolysis of a salt is the action between the ions
of water and the ions of a dissolved salt, forming an acid and a base.
Hydrolysis also explains the basic nature of Na2CO3, thus:
2Na+ + (C0*)~ + 2HOH <=± 2Na+ + 2(OH)~ + H2CO8
Since the H2CO3 is a weak acid, there will be some union of H+
and (CO8) — , thus forming the slightly dissociated H2CO3, and lib-
erating an excess of OH ions which form the strong base, NaOH.
A solution of sodium chloride and water is perfectly neutral, since
it is made from a strong acid and a strong base. A solution of a salt
made from a strong base and a weak acid, or from a weak base and
a strong acid, does not show a neutral reaction.
PRACTICE WORK ON EQUATIONS REPRESENTING THE HYDROLYSIS
OF COMPOUNDS
Complete the following equations, and tell whether the
solution in each case will be acidic, basic, or neutral toward
litmus:
IONS AND DISSOCIATION 241
a) K2CO3 + 2HOH -» d) Na2SO4 + 2HOH -»
b) ZnSO4 + 2HOH-» e) KNO3 + HOH ->
c) A1C13 + 3HOH -> f) K2SO4 + 2HOH -»
What are reversible reactions? Most reactions we deal with are
reactions that take place in water solutions and, hence, are reactions
between ions. The presence of free ions facilitates chemical changes.
If we dissolve sodium chloride, NaCl, and potassium nitrate, KNO3,
in water, these salts immediately dissociate, forming free Na+, Cl~,
K+, and (NO3) ~. These swiftly moving ions constantly meet and
form molecules of KNO3 and NaCl. The entire reaction, therefore,
is said to be reversible, that is, it goes in both directions.
KNO3 + NaCl *± KC1 + NaNO3
A reversible reaction always reaches a point at which change is no
longer apparent. In other words, the reaction has reached a point of
balance or equilibrium. This does not mean that nothing is happen-
ing. On the contrary, the equilibrium is dynamic, or moving, for the
substances are breaking up as rapidly as they are being formed.
Reversible reactions and equilibrium. The dissociation of an elec-
trolyte is a reversible reaction. Thus, when acetic acid, HC2H3O2, is
dissolved in water, free hydrogen ions and free acetate ions are
formed. These ions meet and form acetic acid.
HC2H302 *=± H+ + (CzHsOz)-
Finally a time is reached when the rate of change from free ions
to HC2H3O2 will just equal the rate of change from HC2H3O2 to
free ions. This will happen, in the case of a 0.0 IN acid solution,
when 98.7 percent of the HC2H3O2 is in the form of free ions and 1.3
percent is in the form of HC2H3O2. When this condition is reached, it
would appear that the HC2H3O2 is no longer dissociating and that a
condition of stable equilibrium has been reached. In reality the equi-
librium is not stable but dynamic, or changing. Changes go on even
after the 98.7 : 1.3 ratio is attained, but while new HC2H3O2 is being
formed, more is being ionized, keeping the same ratio.
When do reactions go to completion? Substances that do not
dissociate cannot react reaoily with other substances that do. If even
one of the products of a reaction is unable to dissociate to any great
extent, the backward action cannot take place. The reaction is then
said to go to completion. Advantage is taken of this to secure reac-
tions that complete themselves.
Fig. 52. Action of silver nitrate and dilute hydrochloric acid.
Does this reaction go to completion? Why? Under what
conditions do reactions go to completion?
Insoluble substances, called precipitates, gases liberated under nor-
mal temperature conditions, and pure water are practically incapable
of dissociation. Hence a reaction will go to completion whenever one
of the products formed is (1) a precipitate, (2) a gas, or (2) water.
Examples of such reactions follow:
1) Formation of a precipitate. Chemists use as tests reactions in
which precipitates are formed. Thus, in testing for a chloride, silver
nitrate, AgNO;{, is added. The formation of insoluble silver chloride,
AgCl, prevents the reaction from reversing itself.
Ag+ + (N03)- + Na+
AgCl [ + Na+ + (NO8)-
2) Formation of a gas. IrTthe laboratory preparation of hydrogen
chloride, a mixture of sodium chloride and sulfuric acid is heated,
and the hydrogen chloride gas that is liberated leaves the field of ac-
tion. Hence the reaction does not reverse itself.
Na+
l- + H+ + (HS04) ~ -» HC1
(HSO4)
3) Formation of undissociated water. During neutralization, wa-
ter is always one of the products formed. Therefore the neutralizing
reaction is complete, since water, which is practically undissociated,
may be considered as having left the field of chemical action.
Na+ + (OH)- + H+ -f Cl- -» Na+ -f Cl~ + HOH
This is strictly true only when the salt formed is not hydrolyzed.
PRACTICE WORK ON EQUATIONS REPRESENTING REACTIONS THAT
GO TO COMPLETION
1. Complete and balance the following equations:
a) KNO3 + NaCl -» c) NaOH + PbCl2 ->
b) BaCl2 -f K2SO4 -» d) KOH 4- H2SO4 -»
242
IONS AND DISSOCIATION 243
e) Na2S03 + H2S04 -> SO, + H2O +
f) FeS 4- HC1 ->
g) CaCO3 -f HN03 -> CO2 + H2O +
h) CaCl2 + NaN03 ->
i) Na2SiO3 -f Ca(OH)2 ->
2. Examine the foregoing equations and, in each case, see
whether any of the products formed are gases or precipitates.
Mark them with the appropriate arrows, j or j. Remember
that S02, H2S, and CO, are gases.
3. Finally, with the aid of the following table of solubilities,
tell whether each of the reactions goes to completion or not,
stating your reason in each case.
TABLE 10.
All nitrates, chlorates, and acetates are soluble in water. SOLUBILITY OF
All chlorides, bromides, and iodides are soluble, except those of Ag, Pb, and Hg. JULUBILI T ur
All sul/ates are soluble, except PbS04, BaS04, CaSO,, HgSO,, and Ag,S04. COMMON
All hydroxides are insoluble, except those of Na, K, NH,, Ca, and Ba. COMPOUNDS
All sulfides are insoluble, except those of Na, K, NH4, Ca, and Ba.
All oxides are insoluble, except those of Na, K, and Ca.
All carbonates, silicates, and phosphates are insoluble, except those of Na, K, and NH.,.
The law of mass action and equilibrium, The quantity of a sub-
stance in a unit volume of solution is a measure of the concentration
of a solution. As early as 1803, the French chemist Claude Berthollet
(who first made use of the bleaching action of chlorine) noticed that
the direction of a chemical reaction is dependent upon the concentra-
tions of the substances involved in the reaction.
He noticed that, in general, the greater the concentration, the
greater the speed of the reaction. For example, a match burns quietly
in ordinary air which contains about 21 percent oxygen. In pure oxy-
gen, however, the match burns much more quickly, since the con-
centration of the oxygen (one of the reacting substances) has been
increased almost fivefold. A greater concentration of oxygen means
more molecules of oxygen per unit volume of gas, and, therefore, a
greater possibility for oxygen molecules to come in contact with
molecules of carbon and carbon compounds. This causes an increase
in the speed of the chemical reaction involved in burning.
Sixty-four years later, Guldberg and Waage, professors of mathe-
matics and chemistry at the University of Oslo, Norway, demon-
strated that the speed of a chemical reaction is directly proportional
to the concentrations of the reacting substances. They also made an
interesting discovery concerning the point of dynamic equilibrium
244
NEW WORLD OF CHEMISTRY
of a chemical reaction — that is, the poiixf at which the reaction
in one direction just balances the reaction in the opposite direction.
They found that a chemical reaction which is normally reversible
can be forced to go in one direction with small reversal. This is ac-
complished by manipulating the concentrations of the reacting sub-
stances.
Guldberg and Waage expressed this phenomenon in the form of
the law of mass action, which implies that a change in the quantity
of the reacting substances results in a change in the equilibrium point
of the reaction. In the manufacture of chemicals, the direction of
a reaction is so controlled that large yields are produced.
How the addition of a common ion forces a reaction to go to
completion. In the light of the law of mass action, let us consider
a saturated solution of sodium chloride. We may express the reac-
tion that is taking place as:
NaCl <=> Na+ + Cl~
In such a reaction, the product of the concentrations of the free so-
dium ions and the free chloride ions is a constant. If, by any means,
we increase the number of chloride ions, the number of sodium ions
must decrease. The number of sodium ions can decrease only if some
of the sodium chloride comes out of solution.
To increase the chloride ions, we add to the solution a compound
of chlorine that dissociates to a high degree. Hydrochloric acid is
such a compound. Therefore, if hydrochloric acid is added to a
saturated solution of sodium chloride, some of the sodium chloride
will be precipitated. We call this shifting of the equilibrium point
the common-ion effect. In this case, the common-ion effect is caused
by the addition of the chloride ion, which is common to HCi and
NaCl.
Fig. 53. Apparatus used to
show the common-ion effect.
NaCl + H2SO4
J^ saturated
NaCl solution
/NaCl
precipitated out
Svante Arrhenius (right) and his close friend, Wilhelm Ostwald. Each was awarded
a Nobel prize for chemistry, Arrhenius in 1903, Ostwald in 1909.
Almost insoluble salts, such as silver chloride, AgCl, may be com-
pletely precipitated by adding a common ion. Thus, the addition of
an excess of NaCl increases the concentration of the few dissociated
chloride ions from the silver chloride, and causes some more AgCl
to precipitate.
A weak acid may be weakened by adding a salt of the weak acid.
Thus the addition of sodium citrate to citric acid weakens that acid,
because the addition of the common ion (the citrate ion) forces
more of the citric acid to the undissociated form.
YOU WILL ENJOY READING
Jaffe, Bernard. Crucibles: The Story of Chemistry, pp. 219-
211. Simon and Schuster, New York, 1948. The story oi the life
and work of Arrhenius.
Samrnis, Constance S. "How Annabella Learned the Facts
about pH." Journal of Chemical Education, Oct., 1942, pp. 490-
494. A delightful, cleverly illustrated treatment of a very im-
portant topic.
USEFUL IDEAS DEVELOPED
1. Science must always answer at least two questions —
how? and why?
2. Two great theories underlie much of chemistry: (1) the
electron theory and (2) the theory of dissociation.
245
246 NEW WORLD OF CHEMISTRY
3. The theory of ionization was advanced to explain elec-
trolysis, nonelectrolytes, strong and weak acids and bases, and
other puzzling facts.
4. An ion is an atom or radical that carries an electric
charge.
5. When an acid, base, or salt goes into solution, it dis-
sociates, partially forming free ions. When a current is passed
through such a solution, the ions are attracted to the electrodes
having a charge opposite their own.
6. Free ions are formed in greatest numbers in water so-
lutions of electrolytes. Free ions occur also in molten sub-
stances, in heated gases, and in the air surrounding radio-
active substances. Atoms of glowing metals throw off electrons;
if the electrons come in contact with a gas, as in a radio tube,
the molecules of the gas become ionized.
7. In terms of ionization, an acid is a substance that pro-
duces free hydrogen ions; a base is a substance that produces
free hydroxyl ions.
8. A strong acid is one that dissociates easily, forming large
numbers of hydrogen ions. A weak acid forms only a small
number of hydrogen ions. Strong bases form a large number of
hydroxyl ions; weak bases form few. A compound, such as
water, whose molecule acts like a tiny magnet, is called a
dipole. The molecules of nonpolar compounds, such as CO2,
do not dissociate in solution, are nonelectrolytes and noncon-
ductors of electricity.
9. Neutralization is the union of the hydrogen ions of an
acid with the hydroxyl ions of a base.
10. The acidity or alkalinity of a solution is measured in pH
values. Water, with a pH of seven, represents neutrality. A sub-
stance with a pn greater than seven is basic; one with a pn less
than seven is acidic.
11. Substances dissociate to different degrees, depending
on the (1) solute, (2) solvent, (3) concentration of the solu-
tion, and (4) temperature.
12. All water solutions of acids, bases, and salts are elec-
trolytes; that is, they dissociate and conduct an electric current.
13. Electrolytes raise the boiling points and lower the
freezing points of solutions to a greater degree than non-
electrolytes do, because they dissociate, forming larger num-
bers of ions.
14. The electron theory upholds the theory of dissociation.
The charge on an ion is the same as the valence of the element.
15. Hydrolysis of a salt is the action between the ions of
water and the ions of a dissolved salt, forming an acid and a
base. It is the opposite of neutralization. Salts formed by the
IONS AND DISSOCIATION 247
reaction of a strong acid and a weak base are acidic; by the
reaction of a weak acid and a strong base are basic.
16. A reversible reaction is one that will go in two direc-
tions depending upon the conditions of the reaction. Since
most chemical reactions are reactions between free ions, if the
ions are neutralized or removed, the reaction goes to com-
pletion.
17. Reactions go to completion when one of the products
formed is (1) a precipitate, (2) a gas, or (3) water.
18. The common-ion effect is caused by adding an ion iden-
tical with one of the ions of a compound in solution. It results
in a partial precipitation of that compound.
USING WHAT YOU HAVE LEARNED
Group A
1 . To what questions did Arrhenius seek an answer?
2. (a) What is an electrolyte? (b) What three classes of
compounds are electrolytes? (c) Give three examples of non-
electrolytes.
3. How could you find out whether a solution contained
an electrolyte?
4. Explain in terms of the dissociation theory what hap-
pens when NaCl is dissolved in H^O.
5. Name two ways in which ions differ from atoms.
6. In terms of the theory of dissociation define (a) an acid,
(b) a base.
7. (a) On what do the properties common to bases de-
pend? (b) to acids?
8. Why do we always use dilute acid in preparing hy-
drogen by displacement of hydrogen from that acid by a
metal?
9. Why will a thoroughly insoluble hydroxide not turn
pink litmus blue?
10. Explain neutralization in terms of ions.
1 1 . Define strong acid and strong base in terms of the dis-
sociation theory.
12. What are (a) three common strong acids, (b) two com-
mon weak acids, (c) three common weak bases, and (d) two-
common strong bases?
248 NEW WORLD OF CHEMISTRY
13. A solution has a pn value of six. What does this mean?
14. Normal blood has a pn of about 7.3. Is it acid or alka-
line?
15. How does the degree of dissociation of NH4OH com-
pare with that of NaOH in equivalent solutions?
16. (a) What happens to a metal atom when it becomes an
ion? (b) Explain the dissociation of NaCl in terms of elec-
trons.
17. Explain the difference in physical and chemical prop-
erties of the potassium atom and the potassium ion.
18. By means of diagrams, show the difference between the
chloride ion and the chlorine atom.
19. Change the following equations into ionic equations.
Consult the table of solubilities on page 243.
a) AgNO3 + KC1 -* AgCl + KNO8
b) 2NaCl + H2SO4 -» 2HC1 + Na2SO4
c) BaCl2 + H2S04 -» BaSO4 + 2HC1
20. (a) What is a concentrated acid? (b) Would a concen-
trated acid necessarily be a strong acid? Explain.
21. Insoluble bases are very weak bases. Explain.
22. Complete and balance the following equations.
a) Na2C03 + CaCl2 -> f) Pb(NO3)2 + NaCl -»
b) Zn + H2S04-» g) BaCl2 + K2SO4 ->
c) AgNO8 + NaBr -> h) NaNO3 -f CuCl2 ->
d) NaCl + H2SO4 -> i) Pb(NO3)2 + H2S ->
e) NaCl + Zn(NO3)2 -> j) NaOH + NH4C1 -»
23. Which of the reactions in exercise 22 are reversible and
which go to completion?
24. (a) Which of the following salts are neutral in solution?
(b) Which give an acid reaction? (c) basic reaction? (d) Ex-
plain.
• (1) NaCl 1(3) A1C13 l/(5)Na3BO8 (7) ZnSO4
Na2CO3 (4) KNO3 (6) K2SO4 J£) Na2SO8
25. What is the percentage composition of sodium acetate,
NaC2H302?
IONS AND DISSOCIATION 249
26. Explain the electrolysis of H2O in terms of ions and
electrons.
Group B
27. Why does cold dilute H2SO4 attack Zn, although cold
concentrated H2SO4 does not?
28. Dry cobalt chloride (CoCl2) is blue in color. A solution
of this salt in water is pink. Explain why this color change
occurs.
29. In what three respects has the original theory of disso-
ciation been modified?
30. Concentrated H2SO4 is a poorer conductor of electricity
than dilute H2SO4. What does this indicate with reference to
ions?
31. Explain how the study of the freezing points of solutions
led to the theory of dissociation.
32. (a) State the law of mass action, (b) Explain how it
works, using a suitable reaction as an example.
33. What is a dipole? Illustrate your answer.
34. What weight of "Prestone," C2H4 (OH) 2, should you
use in a 3| gallon auto radiator to protect it against 5°F.
weather? One mole of C2H4 (OH) 2 lowers the freezing point
of one liter of water from 0°C. to -1.87°C. Density of
C2H4(OH)2 is 1.13.
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Write a report or organize a class discussion on the im-
portance of pn measurements to a soil (agricultural) chemist
or to a medical research man. Consult your teacher of agri-
culture or your family doctor.
2. Construct a laboratory setup as shown on page 239. Plate
some copper onto a piece of steel. The steel object should be
the cathode and the copper strip the anode. The electrolyte
should be a copper sulfate solution.
3. Take an old radio tube apart and demonstrate to the
class how electrons are liberated ifrom the heated filament, and
how the flow of electrons passes through the grid toward the
positive plate. Consult your teacher of physics or a radio
engineer in connection with this project.
17
AMMONIA
AND REVERSIBLE REACTIONS
The men of experiment are like
the ant: they only collect and use;
the reasoneis resemble spiders who
make cobwebs out of their own sub-
stance. But the bee takes a middle
course; it gathers its material from
the flowers of the garden but trans-
forms it by a power of its ozvn.
Francis Bacon (1561-1626) in No-
vum'Oiganum
The electron structure of ammonia. At high temperatures, nitro-
gen combines with hydrogen, forming ammonia, NHV Nitrogen acts
like a nonmetal with a valence of three. Since its outer ring is short
three electrons, it combines with three hydrogen atoms, iorming a
nonionic compound. )
The three-dimensional, or cubic, diagram (Fig. 54) is another
method of showing the arrangement of electrons around the nu-
cleus of an atom. This diagram represents the formation ot ammonia.
Note that the electrons of the three hydrogen atoms are shared by
the nitrogen atom in forming the ammonia molecule. Since NH,
bears two unshared electrons, it can accept, or combine with a pro-
ton (H+) to form NH4+. Hence, it is a base.
How ammonia is prepared in the laboratory. When ammonium
chloride is heated with calcium hydroxide, ammonia gas is liber-
ated and may be collected by the displacement of air (Fig. 55) .
2NH4C1 + Ca(OH)2 -> CaCl2 + 2NH3 1 + 2H2O
Since the only function of the ammonium chloride is to furnish the
ammonium group, NH4, almost any other ammonium salt may be
250
Nitrogen
Hydrogen
Ammonia NH3
Fig. 54. Formation of ammonia, showing electron transfer*
substituted for it. Since any soluble base can supply the OH radical,
we can use any soluble base instead of calcium hydroxide. The gen-
eral method for preparing ammonia is, therefore, by the reaction of
an ammonium salt with any soluble base.
Physical properties of ammonia. Priestley collected ammonia over
mercury, because it is soluble in water. At room temperature 1 vol-
ume of water will dissolve more than 700 volumes of this gas, NH.{.
The extreme solubility of ammonia may be shown by the ammonia
fountain. This setup consists of a flask and a tube (Fig. 56) . The
flask is filled with dry ammonia gas, and inverted over water. As
soon as the tube enters the water, some ammonia dissolves in it,
reducing the pressure inside the flask. The air pressure outside then
forces the water up into the tube, and as it issues from the top of
the tube it forms a fountain of ammonia water. Why?
The characteristic pungent odor of ammonia was known to
Priestley. It is reported that as he bent over the fireplace where he
prepared the gas by the method we use today, its vapor made his
eyes fill with tears and drove the occupants of his house out of doors.
Ammonia gas is about half as heavy as air. It may be readily lique-
fied, using only 70 pounds of pressure per square inch at ordinary
temperatures. The colorless liquid NH3 is kept in steel cylinders and
shipped in tank cars.
Fig. 55. Laboratory preparation of ammonia.
Why is the generating tube tilted downward?
Fig. 56. The ammonia fountain.
water
251
252
NEW WORLD OF CHEMISTRY
Chemical properties of ammonia. Ammonia reacts with water,
forming ammonium and hydroxyl ions. The latter ions account for
the basic characteristics shown by the water solution of ammonia,
aqua ammonia.
NH3 + H20 *± (NH4)+ + (OH) -
Ammonia gas does not burn in air, but it does burn in pure oxy-
gen with a pale greenish flame.
4NH3 + 5O2 -» 6H2O + 4NO
Dry ammonia does not unite with dry hydrogen chloride, but the
presence of a trace of water causes the two to combine, forming a
white cloud that settles out as ammonium chloride powder.
NH3 + HC1 <=t NH4C1
Ammonia unites with the very active metals. When passed over
hot magnesium, for example, magnesium nitride and hydrogen are
formed.
2NH3 + 3Mg -> 3H2 + Mg3N2
Ammonia is a very useful compound. Liquefied ammonia gas, or
anhydrous ammonia as it is called, is a widely used compound. In a
recent normal year, more than 2,000,000 tons were produced in this
country for use in the chemical industries.
More than 75 percent of all chemical nitrogen products are pro-
duced from ammonia. Ammonia is the chief raw material from
which nitric acid and nitrates, or salts of nitric acid, are made. Com-
bined with acids to form ammonium sulfate, (NH4)2SO4, ammo-
nium nitrate, NH4NO3, and monoammonium phosphate, NH4H3PO4,
ammonia is one of the chief sources of nitrogen in fertilizers. Urea,
used as a fertilizer, is also made from ammonia. The manufacture
AMMONIA AND REVERSIBLE REACTIONS
253
In recent years, there has been an increasing use of anhydrous
ammonia as a primary fertilizer. Sometimes the ammonia is added
to irrigation water which effectively spreads it throughout the soil
A more widespread practice is to inject liquid anhydrous ammonia
directly into the soil through tubes mounted on a plow-like ap-
plicator. Upon release, the ammonia, which has been kept under
pressure, reverts to a gas and unites chemically with the soil par-
ticles. Greatly increased crop yields have resulted from treating
the soil with anhydrous ammonia.
Ammonia is widely used as a refrigerant. Dry ammonia gas is used
extensively in commercial refrigeration because of the ease with
which it can be liquefied. In making artificial ice, ammonia gas is
placed m a closed system of pipes and coils. By means of a pump,
the gas is compressed until it changes to a liquid; the heat evolved
in the process is removed by a spray of cold water. Then, as cold
hquid ammonia, it is passed through pipes into the freezing cham-
ber, which contains brine or a water solution of calcium chlo-
ride, C_jaCvd9.
As fast as the liquid ammonia enters the pipes in the chamber
through a needle valve, it expands suddenly and vaporizes as a result
ot the reduced pressure, and in so doing absorbs a great amount of
heat from the brine. So cold does the brine become that the pure
water m the tanks that are immersed in the brine changes to blocks
of ice. The ammonia gas returns to the pumping chamber, where it
is again compressed for reuse. The process is continuous, and the
same ammonia gas is used over and over again.
Ammonia gas also helps purify water. One method of water puri-
fication of increasingly wide use employs both chlorine and ammo-
nia. This method is known as the chloramine process and depends
cold water
liquid ammonia-*-
1 -— . •=•-•••:•" "• .-.-.•-^" *•
brine
Fig. 57. Ammonia refrigeration equipment.
on the fact that when ammonia and chlorine react, chloramine,
NH2C1, is formed. One of the reactions is:
2NH, + C12 -> NH2C1 -f NH4C1
Chloramine is a very effective killer of bacteria, or bactericide, but
is less active chemically than chlorine and produces less of the typical
taste of chlorinated water. The process is somewhat less expensive
than chlorination, and is especially well adapted to communities
whose water supplies have musty or swampy tastes.
Preparation and properties of ammonium hydroxide. Care must
be taken to distinguish between ammonia, NH3, and the ammo-
nium ion, NH4+. In the laboratory preparation of ammonia, the ni-
trogen compound first formed might be thought of as ammonium
hydroxide, which breaks down into ammonia and water.
2NH4C1 + Ca(OH)2 -> CaCl2 -f [2(NH4)+ + 2 (OH)-]
I
2NH3 + 2H20
A water solution of NH3 is often called ammonium hydroxide even
though the compound probably does not exist. The solution really
consists of some dissolved NH3, some NH4 ions, and some OH ions. A
saturated solution of ammonia is lighter than water (sp. gr. 0.88) ,
and contains about 36 percent NH8 by weight at room temperatures.
The ammonia may be expelled by boiling. Household ammonia is
a water solution of NH3 containing about 6 percent oleic acid.
A water solution of ammonia is a weak base, since NH3 reacts with
water to form only a few (one percent) hydroxyl ions. Because it is
a base, it dissolves grease and hence removes dirt. Since this ammo-
nia water gives off vapor, or volatilizes, rapidly and completely with-
out leaving a solid, it is useful as a household cleansing agent.
254
AMMONIA AND REVERSIBLE REACTIONS
255
The ammonium radical and ammonium salts. The ammonium
radical, because of its positive valence, is considered metallic. Its
presence may be detected by adding the salt to be tested to a base,
such as sodium hydroxide, and heating the mixture. If the sus-
pected substance is an ammonium compound, ammonia gas will be
liberated, and can be identified by its odor. Moist pink litmus paper
held in the gas turns blue. Why must the litmus be wet?
(NH4)2S04 4 2NaOH -+ Na2SO4 4- 2H2O 4- 2NH3 T
One of the most common of the ammonium salts is ammonium
chloride, or sal ammoniac. This salt, first produced in Egypt, was
known to the early alchemists. It is a white, crystalline substance,
readily soluble in water. It is decomposed by heat into two gases,
NH3 and HC1, which reunite on cooling.
,, NH4C1 *=> NH3 4- HC1
Sal ammoniac is used extensively in dry batteries as an electrolyte
and in soldering in which the hot iron dissociates the salt. The hydro-
gen chloride arid ammonia liberated remove the rust that covers the
surface of the metal to be soldered, the hydrogen chloride by dissolv-
ing the rust, and the ammonia by reducing it.
Fritz Haber makes the synthesis of ammonia a commercial suc-
cess. Minor quantities of ammonia are still made by an old process —
the destructive distillation of coal. This process, in which ammonia
is produced as a valuable byproduct, is discussed on pages 862— 363.
Synthetic methods have superseded the method of preparing am-
monia from coal. The most successful synthesis of ammonia is based
A compressor used in the manufacture of ammonia. The hydogen is made from
natural gas produced in nearby oilfields.
Shell Chemical Corporation
256 NEW WORLD OF CHEMISTRY
on the process first worked out on a commercial basis by Fritz Haber
(ha'ber) in 1913 and known as the Haber process.
The first real test of this great achievement came during World
War I, partly as a result of the desperate need of the German gov-
ernment for nitrogen compounds. Haber's process made agriculture
in blockaded Germany independent of Chile saltpeter and also gave
the German military machine a new source of nitrates for high ex-
plosives. The Haber process enabled Germany to fight hunger, and
stave off defeat much longer than the Allies expected. Fritz Haber
was later forced into exile by the Nazis and died in Switzerland.
The Haber process. This process, with its many modifications, is
the most important single process for the fixation of nitrogen, that is,
the combining of the free nitrogen of the air with other elements to
form useful compounds. The process is based upon what appears to
be a very simple reaction, the union of hydrogen and nitrogen gases.
The nitrogen and hydrogen for this reaction are obtained from
air, coke-oven gas, water gas, natural gas, and some petroleum refin-
ery gases. Although this chemical reaction has been known for a
long time, it was not industrially practicable until the reversible
reaction could be controlled. The laws of chemical equilibrium had
to be used, so that some of the ammonia gas formed would not be
immediately decomposed into its elements.
What factors can be used to control the point of equilibrium?
Most chemical reactions are reversible. In the synthesis of ammonia,
the first equilibrium ratio was about two percent ammonia to 98 per-
cent of a mixture of nitrogen and hydrogen. In other words, most of
the ammonia formed during the union of nitrogen and hydrogen de-
composed into its constituent gases.
The point of equilibrium, however, can be controlled to some
extent. The factors that help to control it are (1) temperature,
(2) pressure, and (3) concentration of the substances involved in
the reaction. Catalytic agents increase the speed of a reaction and,
hence, enable the point of equilibrium to be reached more rapidly.
It is not probable that catalysts alter the point of equilibrium.
Thus, in the case of the preparation of ammonia in America by a
modified Haber process, the reaction
3H2 + N2 -* 2NH3 T + 24,000 calories
has been forced to go to the right, producing as much as 30 percent
NH3, instead of only two percent of the theoretical yield.
AMMONIA AND REVERSIBLE REACTIONS
257
\itrogen Division, Allied Chemicul rind Dye Corporation
Aerial view of a large plant for the synthesis of ammonia and the manufacture
of fertilizer.
The conditions that made the American process a success were
(1) the use of a specially prepared iron oxide as catalyst, (2) a tem-
perature of about 475°C., (3) a reaction pressure of about 300 atmos-
pheres (atrn.) , that is, 300 times atmospheric pressure, and (4) a
rapid removal of the NH3 formed. Since heat is evolved during the
synthesis of ammonia, the higher the temperature, the less the yield,
and consequently too high a temperature is avoided. Since there is
a diminution of volume (three volumes of hydrogen unite with one
volume of nitrogen, forming only two volumes of ammonia) , the
yield is increased by an increase in pressure. The manufacturing
conditions represent the most effective compromise between largest
yield, shortest time, and most profitable rate.
Another synthetic process for making ammonia. In 1916 the
United States built a large nitrogen-fixation plant at Muscle Shoals,
Alabama, where ammonia was to be prepared by the cyanamide proc-
ess. This project was later abandoned, although the cyanamide
process at Niagara Falls, Canada, has been very successful as a source
of calcium cyanamide. The chief chemical changes that take place
in the cyanamide process are:
1) The formation of calcium carbide by heating lime, CaO, and
coke, C, in an electric furnace.
CaO + 3C -» CO -j- CaC2 (calcium carbide)
258 NEW WORLD OF CHEMISTRY
2) The union of calcium carbide with free nitrogen, forming
calcium cyanamide.
CaC2 + N2 — >• C + CaCN2 (calcium cyanamide)
3) The addition of steam to cyanamide, forming ammonia.
CaCN2 + 3H2O -» CaCO8 + 2NH3 f
In dry form, crude, powdered calcium cyanamide containing about
60 percent CaCN2 is sold as a fertilizer under the name of "Cyana-
mid." All of its nitrogen is available as a plant food.
YOU WILL ENJOY READING
Berl, E. "Fritz Haber." Journal of Chemical Education, May,
1937, pp. 203-207. A short biography.
Jaffe, Bernard. Men of Science in America, pp. 307-330.
Simon & Schuster, New York, 1944. The development of our
ideas regarding reversible reactions is tied up with the contri-
bution of America's greatest theoretical scientist, J. Willard
Gibbs. His life and work are described here.
Ross, William H.; Adams, J. Richard; Yee, J. Y.; and Whit-
taker, Colin W. "Preparation of NH4NO3 for Fertilizer Use."
Industrial and Engineering Chemistry, Dec., 1944, pp. 1088-
1095.
Slosson, Edwin E., Creative Chemistry, pp. 14-36. D. Apple-
ton-Century Co., New York, 1920. "Nitrogen, Preserver and
Destroyer ol Lite."
USEFUL IDEAS DEVELOPED
1. When a substance which ionizes goes into solution, the
change that takes place is reversible. The substance dissociates
into ions; these ions unite, re-forming the original substance.
Even when equilibrium is established, this reversible reaction
continues. Change is no longer apparent, however, because
the rate of dissociation is the same as the rate at which the ions
in the solution reunite, forming the original substance. The
solution is in a state of dynamic equilibrium.
2. The point of equilibrium ot a reversible reaction can be
controlled to a certain extent by (1) temperature, (2) pres-
sure, and (3) concentration of the substances involved. Cata-
lytic agents increase the speed of a reaction and, hence, enable
the point of equilibrium to be reached more rapidly.
AMMONIA AND REVERSIBLE REACTIONS 259
USING WHAT YOU HAVE LEARNED
Group A
1. Describe the laboratory preparation of NH3. Use a
labeled diagram.
2. (a) How did Priestley first collect ammonia? (b) Why
did he use this method?
3. (a) What is the general method of preparing NH3?
(b) Write two equations illustrating two ways of preparing
NH3 from NH4C1.
4. Show, by a simple experiment, the extreme solubility of
NH, in H2O.
5. Devise a simple experiment to determine whether NH8
is lighter or heavier than air.
6. Does NH3 burn? Explain.
7. (a) By means of an equation, give the chief chemical
property of NH3. (b) What type of reaction is this? (c) How
can you make the reaction go to the right? (d) to the left?
8. Complete and balance the following equations:
a) NH3 + HC1 -» c) (NH4),S04 + NaOH->
b) NH3 + Mg -> d) NH3 + H2SO4
e)
9. What are five uses of NH3?
10. What substances are present in a water solution of NH3?
11. (a) Is a solution of NH3 in water a strong base? (b) Ex-
plain your answer.
12. (a) Write an ionic equation showing the neutralization
of HC1 by NH4OH, (b) also the neutralization of H2SO4
and (c) of HNO3 by the same base.
13. Does dry NH3 affect litmus? Explain.
14. (a) In what group of elements is the ammonium radical
placed? (b) Explain.
15. What are two commercial methods of preparing NH3?
16. What was the difficulty that confronted manufacturers
who attempted to make NH3 by direct synthesis of its ele-
ments?
17. What four conditions are met in the improved Haber
process?
260 NEW WORLD OF CHEMISTRY
18. (a) Describe the cyanamide process for making NH8.
(b) What are the three equations?
19. Complete and balance the following equations:
a) NH3 + HNO3 -> b) NH3 + H3PO4 ->
20. "Spirits of hartshorn" was the name applied to am-
monia water prepared by the alchemists by heating the horns
of deer. What elements must have been present in deer horns?
21. In what two ways do liquid ammonia and aqua am-
monia differ?
22. What simple test will distinguish NH^Cl from Nad?
Both of these compounds are white, soluble salts.
23. Compare the laboratory methods of collecting NH3, N2,
HC1, 02, C12, and H2.
24. Explain the action of liquid NH3 on dry litmus paper.
25. Why is ammonia water called the volatile alkali?
26. Bottles of household ammonia were formerly closed
with rubber rather than with glass or cork. Why?
27. Why is ammonia water, rather than lye, used in re-
moving grease spots from clothing?
28. Compare the ease with which NH3, N2, C12, and H2 are
liquefied.
29. Describe the preparation of artificial ice by means of
NH3. (b) What property of NH8 makes it useful in this proc-
ess? (c) Of what use is the CaCl0 or brine solution?
'
30. What weight of NH4C1 is necessary to make 340 tons
of dry liquid NH3?
31. Determine the percentage of nitrogen in monoammon-
ium phosphate.
32. (a) What is chloramine? (b) How is it prepared?
Group B
33. Bottles of ammonia water and hydrochloric acid are
placed within a few inches of each other and their stoppers re-
moved. White fumes are seen. Explain.
AMMONIA AND REVERSIBLE REACTIONS 261
34. Explain how you would determine the strength of a bot-
tle of household ammonia by titration. Give details.
35. How can a knowledge of the laws of equilibrium be
used in making the preparation of NH8 by the Haber process
more efficient and more economical?
36. When water is added to magnesium nitride (Mg3N2) ,
the odor of NH3 is detected. Write an equation to explain this.
37. CaCl2 unites with NH3, forming CaCl2 • 8NH3. Explain
whether you could use CaCl2 in drying ammonia gas.
38. Account for the odor of NH3 around heaps of garbage
and manure.
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Take a dry cell apart, and show the various substances
that are used in its manufacture. Prove the presence of NH4C1
by a chemical test. What is the purpose of the NH4C1?
2. Write a two- or three-page report on the freedom of the
man of science in a democracy compared with the enslavement
of science in totalitarian states. Use Haber and, perhaps, Lang-
muir as examples.
3. Set aside 2 ml. of household ammonia in an open test
tube for two weeks. Test for the presence of ammonium hy-
droxide with pink litmus paper at the end of the time. Explain
the results.
18
NITRIC ACID
AND NITROGEN COMPOUNDS
. . . For nitrogen plays a double role
in human economy. It appears like
Brahma in two aspects, Vishnu the
Preserver and Siva the Destroyer.
E. E. Slosson, 1919
The revolution brought about by man-made nitrogen com-
pounds. Haber's successful synthesis of ammonia widened man's
control over nature by making him tess dependent for his raw mate-
rials on supplies present in limited or faraway areas of the earth's
surface. Synthetic ammonia was soon converted into nitric acid.
This, in turn, gave mankind an unlimited supply, not only of ferti-
lizers, but also of high explosives. The synthesis of nitrogen com-
pounds from air and water was a bloodless revolution whose conse-
quences touched the lives of half the people of the world.
Preparation and properties of nitric oxide. Not only does nitrogen
unite with hydrogen at high temperatures, but it combines with oxy-
gen also when an electric spark is passed through a mixture of the
two gases. Cavendish made this discovery in 1770 when he passed an
electric spark through mixtures of hydrogen and air in his synthe-
sis of water. Soon afterward Priestley made a thorough study of the
compound formed by the direct union of nitrogen and oxygen.
N2 + O2 — » 2NO (nitric oxide)
Priestley also prepared nitric oxide by the action of copper on
dilute nitric acid, collecting the gas by the displacement of water,
262
NITRIC ACID AND NITROGEN COMPOUNDS
263
as shown in the illustration below. This is the laboratory method
used today.
3tu + 8HNO3 -> 3Cu(NO3)2 4- 4H2O + 2NO T
Nitric oxide is a poisonous, colorless gas, very slightly soluble in
water, and about as heavy as air. Chemically, it is very active. When
exposed to air or oxygen, it oxidizes at once to nitrogen dioxide, NO2.
This reaction produces much heat; that is, it is decidedly exothermic.
Nitrogen dioxide. The equation representing the oxidation of
nitric oxide to nitrogen dioxide is:
2NO + O2
2NO2
Nitrogen dioxide is reddish brown in color, heavier than air, very
soluble in water, and easily liquefied. Its fumes are irritating and
poisonous.
In preparing nitric oxide by means of copper and nitric acid, the
gas first seen in the generating bottle is brown nitrogen dioxide
rather than the colorless nitric oxide. Actually NO is formed first,
but combines with the oxygen in the air of the generating bottle to
form the brown NCX. When the mixture of the NO and NO2 passes
through the delivery tube into the collecting bottle, the NO2 dis-
solves in the water. Only the colorless nitric oxide displaces the
water in the bottle.
As its temperature is lowered, nitrogen dioxide gradually changes
into nitrogen tetroxide gas, N2O4. The reaction 2NO2 ^ N2O4 is a
reversible one. Above 140°C., the reaction goes to completion, as
shown by the arrow pointing to the left. At room temperature, the
gas obtained is a mixture of the reddish brown NO2 and the color-
less N2O4. At very low temperatures, NO2 changes completely into
N204/
The arc process of making nitric acid from air and water. The
union of nitrogen with oxygen in an electric arc, or by the action of
an electric spark, was used as the basis of many commercial ventures
attempting to produce synthetic nitric acid. In the beginning, they
nitric oxide
f
Fig. 58. Laboratory prepara-
tion of nitric oxide. What other
laboratory setup does this re-
semble?
copper wire
inside glass
^^^^••••••••••••••••o.
Fig. 59. Laboratory preparation of nitrogen dioxide by the arc process. How can
you tell whether NO2 is formed?
all failed because when a mixture of nitrogen and oxygen was thus
treated, the yield of nitric oxide was very small. Because the reac-
tion is reversible, practically all the nitric oxide formed at first was
decomposed into its original elements.
N2 4- O2 ?=± 2NO T
A careful application of the laws of chemical equilibrium finally re-
sulted in larger yields of nitric oxide. This gas combined with the
oxygen of the air to form nitrogen dioxide which was then dissolved
in water to form nitric acid, HNO,.
2NO + 02
3N02 + H20
»2NO2
> 2HN03 4- NO
The Ostwald process, another commercial method of making
nitric acid. The arc process because of its higher cost has been super-
seded by the Ostwald process. No sooner had the synthesis of ammo-
nia been successfully carried out than Wilhelm Ostwald (ost'valt) ,
a chemist who helped Arrhenius establish the theory of ionization,
showed how ammonia gas could be converted into nitric acid on an
industrial scale.
The Ostwald process consists of oxidizing ammonia gas in the pres-
ence of a catalyst, which consists of a specially-prepared platinum or
platinum-rhodium gauze heated to a red heat. The two oxides of
nitrogen are formed during the process, but the final products may
be represented by the following equation:
NH3 + 2O2 -> HN03 + H2O
Using a pressure of only about six atmospheres, about 95 percent
of the ammonia is converted into a water solution of nitric acid of
about 50 percent concentration.
Nitrogen fixation. The change of free nitrogen into ammonia by
the Haber process and the change of free nitrogen from the air into
nitric acid by the arc process are examples of nitrogen fixation.
264
gauze" 1 combustion
exhaust
pymp
H2O
Fig. 60. Laboratory preparation of NOj by the Ostwald process.
Nitrogen fixation is the changing of free nitrogen from the air into
useful compounds. During electric storms, some nitric acid is formed
in the atmosphere by a natural arc process.
Nitrogen-fixing bacteria help the farmer. Most farm crops use
part of the valuable salts in soil. On the other hand, certain crops
such as peas, beans, and alfalfa, actually enrich the soil in which
they are grown. Chemistry explains the fertility of such soil in the
following way: Plants require nitrogen in the form of nitrates.
These nitrates are soluble in water and can, therefore, be absorbed
by the root hairs of plants by osmosis, a process by which liquids and
gases pass through semipermeable tissues. On the roots of plants such
as peas, beans, and alfalfa are little nodules, inside which live bil;
lions of nitrogen-fixing bacteria (Rhizobium) . These bacteria and
several others have the ability to change the free nitrogen, present
in porous soil, into soluble nitrogen compounds, particularly ni-
trates, that plants use in building living tissues. When such crops
are plowed under, this "green manure" adds nitrogen to the soil.
The nitrogen cycle — nitrogen compounds break down into free
nitrogen. Various other bacteria break down nitrogen compounds in
the soil into simpler compounds and even return considerable quan-
tities of free nitrogen to the air. Such bacteria, called denitrifying
(de-ni'trMI-ing) bacteria, cause the loss of nitrogen from the soil and
thus complete the extremely important nitrogen cycle. These bacteria
are responsible, in part, for the rapid decay of nitrogen-containing
These pea roots are covered
with nodules containing nitro-
gen-fixing bacteria.
266 NEW WORLD OF CHEMISTRY
organic wastes. They are used widely in the treatment of sewage (see
pages 222-223) . This series of changes is referred to as the nitrogen
cycle.
How nitric acid is prepared in the laboratory. The laboratory
preparation of nitric acid follows the general method for preparing
an acid. Sodium nitrate, mixed with concentrated sulfuric acid, is
heated gently in a glass retort. Nitric acid is formed, which, having
a lower boiling point than sulfuric acid, evaporates. It is then con-
densed into a colorless liquid by cooling, as shown in the illustration
below. The equation for the preparation of nitric acid is:
NaNO3 + H2SO4 -> HNO3 + NaHSO4 (sodium hydrogen sulfate)
This method is also one of the commercial processes used today in
making about 10 percent of the nitric acid consumed by the world's
industries. Synthetic NaNO3 supplies the nitrate.
Nitric acid was known to the alchemists more than 1000 years ago.
Geber (ga'ber) , an Arabian physician and alchemist, prepared it
about A.D. 800. It was called aqua fortis, meaning strong water.
Physical properties of nitric acid. When pure, nitric acid is a color-
less liquid. Its water solution, containing 68 percent nitric acid by
weight (that is, the concentrated nitric acid of commerce) , has a
specific gravity of 1.4, and boils at 120°C. The concentrated acid
fumes strongly. The yellowish appearance of the nitric acid prepared
in the laboratory is caused by the presence of nitrogen dioxide,
formed by the partial decomposition of the nitric acid during the
heating.
Chemical properties of nitric acid. Because nitric acid mixes with
water in all proportions and dissociates almost completely, thus pro-
ducing large quantities of hydrogen ions, it is a strong acid. Nitric
acid is unstable. In sunlight or when heated, it decomposes into
water, oxygen, and nitrogen dioxide.
4HN03 --» 2H20 + 4NO2 T + O2 T
Fig. 61. Laboratory preparation of nitric acid. By what process is the acid collected?
cold water
nitric acid
NITRIC ACID AND NITROGEN COMPOUNDS 267
A glowing splint inserted in the vapors of boiling concentrated
nitric acid catches fire, thus showing that oxygen is present in the
vapors. When an element, such as nitrogen in HNO3, is in a very
highly oxidized state, that is, it has a high positive charge, the com-
pound is a strong oxidizing agent.
The action of nitric acid on metals illustrates its oxidizing pow-
ers. When hydrochloric acid reacts with many of the metals, hydro-
gen is liberated even if concentrated acid is used. On the other hand,
dilute nitric acid acts on a metal, forming water instead of hydrogen.
Nitric oxide gas is also formed. In fact, one of the methods used to
prepare nitric oxide depends on this action of nitric acid.
3Cu + 8HNO3 -> 3Cu(NO3)2 + 4H2O + 2NO j
When concentrated nitric acid reacts with a metal, nitrogen
dioxide, instead of nitric oxide, is formed. Brown NO2 is produced
by the oxidation of NO to NO..
Cu + 4HN03 -» Cu(N03)2 + 2H2O + 2NO2 1
The electron theory explains the oxidizing power of nitric acid as
follows: oxidation is a loss of electrons; therefore, a substance such
as chlorine, which borrows electrons, is a good oxidizing agent.
HNO3 may be thought of as containing N+++++ , which borrows elec-
trons from the copper in the reaction above, changing to N++ (in
NO) . Thus the copper is said to be oxidized and the nitrogen re-
duced.
Nitric acid attacks proteins, forming a yellow compound. The
yellow coloration produced on the skin by nitric acid is caused by
this chemical reaction (see also Chapter 36) . Nitric acid oxidizes
both cotton and wool, destroying most of the strength of the fibers.
Aqua regia, the acid mixture that dissolves gold and platinum.
Nitric acid, when mixed with hydrochloric acid, oxidizes the latter,
liberating atomic chlorine. This mixture of nitric and hydrochloric
acids as in the equation below is called aqua regia, or royal water.
3HC1 + HNO3 -» 2H2O + NO + 3C1
This chlorine reacts with gold, forming soluble gold chloride.
Au + 3C1 -> AuCl3
Both gold and platinum are insoluble in any one of the three com-
mon strong acids, but they are soluble in aqua regia.
268
FeSO4 + nitrate solution • -j^nar i
jjjji&r concentrated
brown ring--"^^ H2SO4
Fig. 62. The brown-ring test for a nitrate.
How do we test for the nitrate ion? The nitrates, salts of nitric
acid, are all soluble in water, are decomposed by heat, and may
be detected by the brown-ring test. This test is made by adding a
small amount of freshly prepared ferrous sulfate solution to the solu-
tion suspected of containing a nitrate. Concentrated sulfuric acid is
then carefully poured down the side of the test tube in such a way
that it collects at the bottom without mixing with the solution. If
a brown layer forms between the heavy sulfuric acid and the solu-
tion floating on top, a nitrate is present. The sulfuric acid pro-
duces nitric acid by reacting with the nitrate, and the nitric acid
in turn reacts with the ferrous sulfate, forming a brown compound,
FeSO4 • NO.
Gunpowder, the earliest explosive. The inactive element nitrogen
does not unite easily with other elements. And after it does, the
unions so produced are very unstable. In fact, such unions are so
extremely unstable that on the slightest provocation * the nitrogen
breaks away with a bang! Most explosives, except those based on
nuclear fission, depend upon this fact.
Nearly all nonfission explosives contain either nitrate (NO3) , or
nitro (NO2) , radicals. In addition, some explosives contain ammo-
nium radicals, NH4. When compounds that contain nitrate or nitro
radicals are mixed with other compounds that can easily use the oxy-
gen of these unstable radicals, an explosive is the result. In certain
cases, the compound containing the nitrate or nitro radicals actu-
ally supplies the means of its own destruction by furnishing the ele-
ments that can use the oxygen readily. When something, such as a
shock, starts the reaction, the unstable nitrate or nitro radicals re-
lease their oxygen for combination with other elements and liberate
free nitrogen gas. Nearly always, most of the other products of the
reaction are gases also, and because of the high temperatures pro-
duced, terrific pressures result.
The earliest explosive was made thousands of years ago by
the Chinese. How the black powder came to be produced is not
known, but it was made then about as it is now, by mixing approxi-
mately 15 parts by weight of potassium nitrate, KNO3, with three
parts of sulfur, and two parts of powdered charcoal. The resulting
NITRIC ACID AND NITROGEN COMPOUNDS 269
black powder, or gunpowder, explodes with terrific violence, much
lire, and great quantities of acrid smoke. The reaction produced is:
2KNO3 + 3C + S -» K2S + Ns t + 3CO2 T
The Chinese used gunpowder in producing firecrackers and other
kinds of fireworks for use in ceremonies and celebrations. But some-
time within the past few centuries, no one knows for sure just when,
gunpowder was turned to military uses, and, like so many other sci-
entific discoveries, soon created a revolution. The foot soldier with
his primitive musket quickly replaced the heavily-armored and be-
pl umed knight with his lance and shield. Turrets and thick stone
walls were no defense against powerful cannons, and the picturesque
castle of the Middle Ages became obsolete.
Nitrocellulose, nitroglycerin, and some detonators. Today, gun-
powder is considered a relatively "tame" explosive. Since its inven-
tion, chemists have produced many kinds of explosives chiefly by
nitrating, or adding NCX or NO, radicals to, such substances as cot-
ton, glycerin, sugar, starch, and other organic compounds. Nitrocel-
lulose, or guncotton, was produced by Schonbein (shun'bln) in 1846
by nitrating cotton with a mixture of nitric and su If uric acids. Nitro-
glycerin, made by nitrating glycerin, a common byproduct of the
manufacture of soap, was produced in 1847. Both of these compounds
are more powerful explosives than gunpowder, but both are much
more sensitive to shock and, hence, explode much more easily.
In 1888, Alfred Nobel, a Swedish inventor, produced dynamite
by absorbing nitroglycerin in a fine clay, or diatomaceous earth (see
page 500) . Dynamite is much less sensitive to shock and, hence,
Standard Oil Company (\'.J.)
A workman inserts a dynam
cartridge into a drilled h
during the construction of
underground pipeline.
270
NEW WORLD OF CHEMISTRY
can be used with greater safety than nitroglycerin. American dy-
namite usually consists of nitroglycerin absorbed in a wood meal
that resembles fine sawdust. The substance is packed in sticks com-
posed of parchment paper. So overcome was Nobel by the possible
uses of his invention that he dedicated part of its profits to the es-
tablishment of the Nobel Peace prize for outstanding contributions
to the peace of the nations of the world, and of the Nobel prizes for
outstanding contributions to research in physics, chemistry, medi-
cine and physiology, and literature. Peacetime uses of explosives in-
clude mining, building dams, and other construction work.
Explosives are set off, or detonated, by means of a shock produced
by even more unstable, and thus more sensitive, compounds of ni-
trogen called detonators. Fulminate of mercury, Hg (CNO) 2, a
widely used detonator, is 1 1 times as sensitive to shock as trinitro-
toluene, or TNT, and twice as sensitive as nitrocellulose. Lead azide,
Pb (N3) 2, another detonator, is half as sensitive to shock as fulminate
of mercury. Such substances are used in making the caps and other
devices with which explosives are set off.
Nitric acid has many peacetime uses. Aside from the production
of explosives, a major industry even in peacetime, the chief use of
nitric acid is in the production of nitrates of organic compounds,
such as nitrocellulose and nitrobenzene. Nitrocellulose is used in
making some photographic films, and many quick-drying lacquers
and enamels, especially for the automobile industry. Nitrocellulose
is used also in the production of many kinds of artificial leathers.
Nitrobenzene is the basic raw material of the aniline, or coal-tar
dye, industry.
Nitric acid also furnishes the oxides of nitrogen required in the
chamber process for the manufacture of sulfuric acid (see page 311).
Nitrates for fertilizers and metallic nitrates are made from nitric
acid. Sodium nitrate and ammonium nitrate are the chief fertilizers
produced. Strontium nitrate is one of the chemicals used in pyro-
technics, the production of fireworks, which consist mainly of flares
NITRIC ACID
Use by approximate percentage
NITRIC ACID AND NITROGEN COMPOUNDS 271
and shells that give off flames and smokes of various colors. A one
percent silver nitrate solution is put in the eyes of newborn babies
to prevent infection that may lead to blindness.
Nitric acid plays an important role in the pickling of steel, in the
etching of engravers' plates for printing, and in the manufacture of
the arsenate insecticides (chiefly lead and calcium arsenates) so
widely used against the boll weevil, and in the spraying of fruit
trees (see page 453) .
The nitrogen situation today represents a chemical revolution.
It should be apparent already that the production of nitrogen com-
pounds is a basic industry. The world's normal consumption of man-
ufactured nitrogen compounds is many millions of tons annually.
At one time, sodium nitrate from Chile and nitrogen compounds ob-
tained as byproducts from the coal industry and slaughterhouses
were the only sources of nitrogen compounds. Today, the total an-
nual amount of fixed nitrogen produced by chemical methods is
many times as great as the total annual consumption of both Chilean
nitrate and all byproduct nitrogen taken together. The story of the
nitrogen industry bears testimony to the widespread development
of synthetic chemistry. It has changed the economic life of millions.
There are six oxides of nitrogen. In addition to the two oxides
of nitrogen already discussed, four other oxides are known. The com-
plete list is: nitrous oxide, N2O; nitric oxide, NO; nitrogen diox-
ide, NO2; nitrogen trioxide, N2O^ nitrogen tetroxide, N.X)4, a pow-
erful oxidizing agent; and nitrogen pentoxide, N2Or,. They illustrate
the law ot multiple proportions and the fact that nitrogen has sev-
eral different valences. Why?
When nitrogen trioxide, N2Oy, is added to water, nitrous acid, a
very unstable acid, is formed.
N2O3 + H2O -> 2HNO2
Similarly, the addition of water to nitrogen pentoxide, N2O5, pro-
duces nitric acid.
N205 + H20 -> 2HN03
These two gases may therefore be said to be the anhydrides (with-
out water) of nitrous and nitric acids, respectively. An acid anhy-
dride is an oxide whose water solution is an acid.
Nitrous oxide, or laughing gas. Priestley was the first to produce
nitrous oxide, N2O, a colorless, heavy gas, slightly sweetish in odor
and somewhat soluble in water. This was about two years before he
A modern hospital operating
room. Suspended from the
table in the left foreground
are tanks of pure oxygen,
carbon dioxide, nitrous oxide
and other gases ready for im-
mediate administration when
needed.
Ohio Chemical and Surgical hqu,
discovered oxygen. He learned that it supported the burning of a
candle better than did ordinary air. It decomposes rather easily into
oxygen and nitrogen. Just before the close of the eighteenth century,
Humphry Davy achieved fame overnight by his discovery of the
physiological effects of this gas. He breathed four gallons of it and
noticed its power to produce a peculiar intoxication, which included
laughing. The poet Samuel Coleridge, as well as other distinguished
persons, came to Davy's London laboratory to experience the thrill
of inhaling this gas. Nitrous oxide is still prepared as Davy made it,
by heating ammonium nitrate.
NH4N03 -> 2H20 + N2O
In 1842 ether was used as the first anesthetic in surgery by Dr.
Crawford W. Long, a country doctor of Georgia. William Morton's
use of ether at the Massachusetts General Hospital in 1846 intro-
duced this anesthetic to the medical world. Two years earlier, Dr.
Horace Wells, a dentist of Hartford, Connecticut, had one of his
teeth extracted after he had anesthetized himself with nitrous oxide.
Today nitrous oxide is still used as an anesthetic in many opera-
tions, especially those of dentistry. It is usually mixed with about
25 percent oxygen and, in cases of more serious operations, with
ether. This mixture of nitrous oxide and oxygen can be breathed
for a considerable period without harmful effects on the circulatory
272
NITRIC ACID AND NITROGEN COMPOUNDS 273
system or on vital organs. Small amounts of nitrous oxide are used
in preserving perishable foods and liquids. Easily liquefied, it is sold
in cylinders. It is used to eject whipped cream at soda fountains.
YOU WILL ENJOY READING
Conant, James Bryant. The Overthrow of the Phlogiston
Theory. Case 2 of the Harvard Case Histories in Experimental
Science. Harvard University Press, Cambridge, Mass., 1950.
Gives an excellent account of Priestley's confusion between
oxygen and nitrous oxide.
Haynes, William. This Chemical Age, pp. 78-94. Alfred A.
Knopf, New York, 1942. A discussion on explosives and their
relation to the dye industry, entitled "Mars: Chemical Dic-
tator."
Slosson, Edwin E. Creative Chemistry, pp. 37-59. D. Apple-
ton-Century Co., New York, 1920. A very readable account of
nitrogen and nitrogen compounds in relation to plants.
USING WHAT YOU HAVE LEARNED
Group A
1. Explain the arc process of making HNO3.
2. Write equations showing two methods of preparing NO.
3. State one chemical and three physical properties of NO.
4. Make a diagram showing the laboratory preparation
of NO.
5. (a) What happens when NO comes in contact with
air? (b) Explain.
6. Under what conditions does NO0 change into N.,O4?
'
7. Write the equation that is the basis of the Ostwald proc-
ess for the synthesis of HNO3.
8. (a) What natural phenomenon results in the formation
of certain oxides of nitrogen? (b) Explain.
9. (a) What is nitrogen fixation? (b) Illustrate.
10. Make a diagram showing the laboratory preparation of
HN08.
11. When a mixture of NaNO3 and H2SO4 is heated in a
retort, HNO3 is formed a little at a time, (a) What four
274 NEW WORLD OF CHEMISTRY
substances are in the retort? (b) Which is removed by heat?
Why? (c) Why do the other substances remain?
12. (a) What property of HNO3 makes it possible to pre-
pare the acid by the laboratory method? (b) Why could not
HC1 be used instead of H2SO4?
13. Write the equation for the laboratory preparation of
HNO3.
14. State four chemical properties of HNO3.
15. (a) How is aqua regia prepared? (b) Its power to dis-
solve gold results from what property? (c) What property of
HNO3 is shown?
16. (a) Why does HNO3 appear to be yellow when prepared
in the laboratory? (b) What property of HNOrf does this color
indicate?
17. (a) What oxide of nitrogen is always formed when
HNO3 decomposes? (b) What oxide of nitrogen is always
formed when HNO3 acts as an oxidizing agent?
18. How would you test for the presence of the nitrate ion?
19. What are the three principal uses of HNO/
20. (a) Most explosives are based upon what fact? (b) What
is a detonator?
21. (a) What is gunpowder? (b) dynamite?
22. Copy and complete the following: The six oxides of
nitrogen illustrate the law of .... The anhydride of HNO3
is . . . , and the anhydride of HNO2 is .... Another name for
N2O is .... N2O was first used as ... by Dr. Horace Wells.
N2O is prepared by heating ....
23. 120 g. of NO are obtained by the action of Cu on
HNO3. How many grams of Cu (NO3) 2 are formed?
24. How do you explain the fact that the reaction of
N2 -(- O2 — » 2NO, although known for more than a century,
could not be used in the preparation of HNO. until rather
recently?
I ...
25. Why is N2O a better supporter of combustion than NO?
26. (a) How does nature restore some of the nitrogen com-
pounds taken from the soil by growing crops? (b) What is
"green manure"?
NITRIC ACID AND NITROGEN COMPOUNDS 275
27. In 1898, Sir William Crookes, one of England's most
eminent chemists, predicted a future world famine caused by
exhaustion of nitrogen compound fertilizers. Why has his pre-
diction not materialized?
28. Make a table showing the properties of the six oxides of
nitrogen.
29. How would you separate Au from Cu in a copper-gold
mixture?
Group B
30. Why cannot HNO3 be used in preparing H2? Answer
this question in the light of the electron theory.
31. (a) Using the electron theory, explain how HNO3 oxi-
dizes Cu. By inspection of the equation, state how many atoms
of Cu are oxidized, (b) How many atoms of nitrogen are re-
duced? (c) How does the total number of electrons lost com-
pare with the total number gained?
32. How could you determine experimentally whether a
gas contained a high percentage of N2O4 and a small amount
of NO2, or was composed almost entirely of NO,?
33. There is less oxygen in nitrous oxide than in nitric
oxide, (a) Which would support burning better? (b) Explain.
34. Nitrates are unstable in the presence of heat. What
product would you expect Cu (NO3) 2 to yield when heated?
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. At Muscle Shoals, Alabama, there is a huge plant for
generating electric power. This plant is capable of producing
thousands of tons of nitrogen compounds a year by fixation.
The U.S. government spent many millions of dollars on this
project. What is the present state of this plant? Consult your
teacher of history or economics or write to your Congressman,
Senator, or the Tennessee Valley Authority at Knoxville, Ten-
nessee. What do you think should be done? Prepare a debate
or discussion of this question.
2. Write a report comparing the changes brought about by
the introduction of gunpowder into warfare with the changes
that the A- and H-bombs may produce.
3. If available at this time, obtain a fresh sample of some
legume such as peas, beans, alfalfa, or soybeans. Examine its
roots for nodules of nitrogen-fixing bacteria. Exhibit to class
and explain its function.
19
MOLECULES:
AVOGADRO'S HYPOTHESIS
. . . In 1858 the atomic theory of
Da I ton was just 50 years old. Stu-
dents at this time were generally un-
familiar with the word molecule, for
chemists spoke as complacently
about an atom of water as about an
atom of oxygen. For the most part,
also, they had never heard of Amedeo
Avogadro. William Tilden, 1921
The battle over the molecules of Avogadro. In 1860, chemical
science was in a turmoil caused by a misunderstanding of the terms
atom and molecule. Chemists had spoken of "atoms of water," which
is a compound, in the same way in which they mentioned atoms of
hydrogen, which is an element. Some used the term compound
atoms. So great was the confusion that finally a congress of chem-
ists was called at Karlsruhe to decide when to use the term atom and
when to use the term molecule, which up to that time had been used
interchangeably.
Among the brilliant men who gathered at Karlsruhe from all parts
of the world was a bearded young Italian, Cannizzaro (kan-net-
sa'ro) . He had come to champion the use of the term molecule in
the sense that it was used in 1811 by Avogadro (a-vo-ga'dro) , a mod-
est professor of chemistry. According to Avogadro, a molecule is the
smallest part of either an element or a compound zvhich has the prop-
erties of that substance.
Avogadro had reached this new meaning of a molecule from his
study of the behavior of gases. He had died four years before the con-
gress was held, but Cannizzaro championed his ideas so successfully
that they were finally accepted by the congress.
276
MOLECULES
277
What Boyle and Charles discovered about the behavior of gases.
In order to understand the significance of Avogadro's contribution
to chemistry, it is necessary to trace the story of the study of gases
after 1660. In that year Robert Boyle, revered by Englishmen as
the father of modern chemistry, discovered that if the pressure on
a gas is doubled — for example, increased from 15 pounds to 30
pounds per square inch — its volume is decreased one-half. Further-
more, he found that this relationship between pressure and volume
does not depend upon the nature of the gas; it is true for all gases.
Thus we have Boyle's law: The volume of a gas varies inversely as
the pressure exerted upon it if the temperature remains constant.
In 1785, Charles, a French scientist, noted that under constant
pressure the volume of a gas increases ^^ of its volume at 0°C. for
each centigrade degree of rise in temperature; in other words, if the
absolute temperature (see page 644) of a gas increases from 273° to
546° (2 X 273°) its volume doubles also. This he found true of any
gas. Thus we have Charles' law: The volume of a gas varies directly
as the absolute (A) temperature, if the pressure on the gas remains
constant.
Although gases act according to Boyle's and Charles' laws under
ordinary temperatures, they do not do so at very high pressures or
very low temperatures. However, a discussion of the way in which
gases act at low temperatures and under high pressures is beyond the
scope of an introductory course in chemistry.
Gay-Lussac finds the law of combining volumes of gases. The
next important advance came 23 years later. Gay-Lussac, the French-
man who collected air samples in a balloon high over Paris, had
long been interested in the study of gas volumes. He knew that one
volume of nitrogen unites with one volume of oxygen, forming two
volumes of nitric oxide. Besides, when he repeated the experiments
of Cavendish and Lavoisier, he found that two volumes ot hydrogen
unite with one volume of oxygen, forming two volumes of water va-
por. This was not a new discovery, but Gay-Lussac suspected that
"other gases might also combine in simple ratios."
plunger I
pressurj
double
air
Fig. 63. A demonstration of
the principles of Boyle's
law.
278 NEW WORLD OF CHEMISTRY
Resuming his researches, Gay-Lussac discovered that one .volume
of hydrogen chloride gas when brought in contact with one volume
of ammonia gas yielded a white powder, with no residue of either
gas. The two gases had combined volume for volume. Furthermore,
he had read that one volume of nitrogen combined with exactly
three volumes of hydrogen, forming exactly two volumes of ammo-
nia gas. This was an arithmetical simplicity of remarkable signifi-
cance. Fractions of volumes of gases were not involved.
1 vol. nitrogen + 1 vol. oxygen —-» 2 vol. nitric oxide
1 vol. oxygen + 2 vol. hydrogen — » 2 vol. water vapor
1 vol. HC1 + 1 vol. NH3 -» NH4C1 (a solid)
1 vol. nitrogen -f 3 vol. hydrogen — * 2 vol. ammonia gas
On the last day of the year 1808, Gay-Lussac formulated from these
observations a law, which bears his name. Gay-Lussac's law states
that the relation between the combining volumes of gases and the
volumes of their products (if they, too, are gases) may be expressed
in small whole numbers. Why this regularity? On the basis of Dai-
ton's atomic theory, chemists could not explain this law.
When Dal ton was faced with this fact, he refused to accept Gay-
Lussac's law. "The truth is," Dalton maintained, "that gases do not
unite in equal or exact proportions in any one instance. When they
appear to do so, it is owing to the inaccuracy of our experiments."
Later, however, after further experimentation and study, Dalton
accepted Gay-Lussac's law. His acceptance of Gay-Lussac's law, when
experimental evidence pointed toward its accuracy, reveals Dalton
as a true scientist.
Is the sum of 2 and 1 always 3? The law of Gay-Lussac in par-
ticular, and the laws of Boyle and Charles to a lesser degree, sug-
gested a number of interesting problems to Avogadro's inquiring
mind. Why, for example, is the behavior of gases so uniform under
changing temperature, while the behavior of solids and liquids is
so variable? Why do gases combine in simple ratios by volume? Fur-
thermore, why, with respect to gases, is not the sum of two and one
always three? For example, why do two volumes of hydrogen unite
with one volume of oxygen, making two volumes of water vapor; and,
similarly, three volumes of hydrogen combine with one volume of
nitrogen, making two volumes of ammonia gas?
Avogadro continues his work with gases. Avogadro tried to answer
these questions. In reference to the last question, perhaps he thought:
it might be that equal volumes of gases contain the same number of
molecules.
MOLECULES 279
But according to this idea, one volume of oxygen ought to combine
with one volume of nitrogen, forming one volume of nitric oxide:
N + O -» NO
whereas, according to actual experiment, two volumes of nitric oxide
are formed. Something was wrong. But wait! Suppose the molecule
of nitrogen gas contains two atoms instead of one, that is, is N2 and
not N, and similarly the oxygen molecule is O2 and not O, what then?
According to this idea, the equation would be:
N2 + O2 -> 2NO
and the conditions necessary for the formation of two volumes of NO
would be fulfilled.
How would this idea work out in other cases, for example, in the
formation of water? If the molecule of hydrogen contains two atoms,
like the molecule of oxygen, we should have the equation,
2H2 + O2 -> 2H2O
or stated in other words, two volumes of hydrogen unite with 1 vol-
ume of oxygen, forming two volumes of water vapor. This agrees
with actual measurements of the volumes.
Apparently Avogadro was on the right track. It remained for him
only to test his hypothesis further by means of other gas combina-
tions to be able to show that, assuming the molecules of elementary
gases to be composed of two atoms each, the volumes corresponded
with the equation as he had calculated.
This he actually did and finally was able to establish the accuracy
of his hypothesis, that all gases behave alike, because equal volumes
of all gases under the same conditions of temperature and pressure
are composed of the same number of molecules (Avogadro's hypoth-
esis) . To this professor from Turin, elementary gases, such as hydro-
gen, oxygen, nitrogen, and chlorine, consist normally of molecules
each composed of two atoms instead of one, as Dalton and the rest
of the world had supposed. Incidentally, Avogadro's hypothesis rec-
onciled the atomic theory of Dalton with Gay-Lussac's law. How?
Of what value to chemistry was Avogadro's hypothesis? What
evidence did Avogadro have to back up so bold an hypothesis? He
could not verify it experimentally. No balance was sensitive enough
to weigh a molecule. It would take billions of these tiny particles
to turn the scales of even the most sensitive balance. He surely had
not looked into the molecules of matter and detected the twin
280 NEW WORLD OF CHEMISTRY
arrangement of atoms, for it would take many millions of mole-
cules placed side by side to make a line one inch long. Only in re-
cent years have methods been developed which can make such a
tiny particle visible.
The only evidence Avogadro had was that of clear, accurate rea-
soning and his own creative imagination. However, this evidence
was strong enough to clear the air and allow chemistry to advance.
The particles of elementary gases were henceforth considered to be
diatomic, that is, composed of two atoms to the molecule. (Later,
by other methods, the inert gases of the atmosphere were shown to
contain only one atom to the molecule.) Atomic weights and molecu-
lar weights were thus clearly differentiated. New methods were made
possible for determining the molecular weights of gases and, from
these, their atomic weights also.
How Avogadro's hypothesis was actually verified. Since the time
of Avogadro, new apparatus and new methods have been devised for
verifying his prophetic statement. A number of scientists, Millikan
and Perrin (pe-raN') among them, determined by experiment the
number of molecules in a given volume of gas. They found that the
number of molecules in two grams of hydrogen gas (a gram-molecu-
lar weight) , for example, is approximately 602,000,000,000,000,000,-
000,000 (602 sextillion) . This number, usually written 6.02 X 1023,
is now called Avogadro's number. Approximately this number of
molecules is known to be present in equal volumes (22.4 liters) of
all gases and vapors under the same conditions of temperature and
pressure. This is no idle guess.
Perrin and Millikan, both Nobel prize winners in physics, main-
tained that we can count the number of molecules in a small volume
of a gas with as much accuracy as we can determine the population
of a city such as New York. Avogadro's hypothesis has now taken its
place as one of the laws of chemistry.
Then came another remarkable verification. Irving Langmuir
(lang'mur) , another Nobel laureate in chemistry, succeeded in break-
ing up the molecules of hydrogen gas. As a result of his experiments
he found that hydrogen gas is made up of molecules each of which
consists of two atoms. Langmuir made use of this discovery when he
invented the atomic-hydrogen torch.
Principle of the atomic-hydrogen torch. In an atomic-hydrogen
torch, hydrogen gas is passed through an electric arc produced be-
tween electrodes made of wolfram. The heat of the electric arc splits
the hydrogen molecule into hydrogen atoms. Immediately after pass-
ing through the arc, the atoms reunite, forming hydrogen molecules,
General Electric Company
Repairing worn parts of vacuum cleaners with an atomic-hydrogen arc welder.
which are oxidized, forming water. Atomic hydrogen cannot be
stored. ^
H2 *=> H + H
All the energy absorbed from the electric arc in splitting the mole-
cule is liberated when the atoms of hydrogen reunite. This energy,
as heat, added to the heat normally generated when hydrogen burns,
produces a temperature between 4000°C. and 5000°C. (The oxy-
acetylene torch gives a temperature of about 3300 °C.)
The atomic-hydrogen torch is used for cutting and welding metals.
It has the advantage of protecting the object against oxidation, since
the jet of burning hydrogen is always surrounded by hydrogen, a
reducing agent.
Proof that the molecule of hydrogen contains two atoms. The-
oretically, we can prove the formula of hydrogen gas to be H, as
follows: (1) From experiments, we know that one volume of hydro-
gen unites with one volume of chlorine, yielding two volumes of hy-
drogen chloride gas. (2) According to Avogadro's law, equal vol-
umes of all gases contain the same number of molecules. Conversely,
equal numbers of molecules of gases occupy equal volumes. Hence,
281
282 NEW WORLD OF CHEMISTRY
one molecule of hydrogen and one molecule of chlorine occupy ,equal
volumes, and two molecules of HC1 gas occupy twice this volume.
We may represent this graphically as follows:
1 volume 4- 1 volume +~ 2 volumes (by experiment)
1 molecule 4- 1 molecule ^ \2 molecules/- --(by Avogadro's
of hydrogen \ of HCI / law)
must contain 1 must contain
' +~ at leasF*2"atoms of hydrogen
(3) One of the HCI molecules must contain at least one atom of hy-
drogen, since fractions of atoms do not exist. (4) Since we have two
HCI molecules, we must have at least two atoms of hydrogen which
can have come from only the one molecule of hydrogen.
PRACTICE PROBLEMS BASED ON THE PROOF OF THE STRUCTURE
OF A MOLECULE OF HYDROGEN
1. We know that 1 vol. of nitrogen unites with 1 vol. of
oxygen, forming 2 vol. of nitric oxide. Prove that the mole-
cule of nitrogen contains two atoms.
2. One vol. of hydrogen unites with 1 vol. of bromine vapor,
forming 2 vol. of hydrogen bromide gas. Show that the mole-
cule of bromine vapor contains two atoms.
3. One vol. of oxygen unites with 2 vol. of hydrogen, form-
ing 2 vol. of water vapor. Show that the formula for oxygen
gas is O2.
4. Prove the structure of the hydrogen molecule from the
fact that 1 vol. of nitrogen unites with 3 vol. of hydrogen,
forming 2 vol. of ammonia gas.
The gram-molecular volume of a gas or vapor. You have learned
that a chemical formula may represent one molecule of a compound;
one molecular weight of a compound; and also one gram-molecular
weight, or mole, ot a compound. For example, CO2 may stand for
one molecule of carbon dioxide, for the molecular weight of car-
bon dioxide (44) , or for one mole of carbon dioxide (44 grams) . If
a coefficient appears in front of a formula, it represents a definite
number of molecules, molecular weights, or moles. Thus 2CO2 stands
for two molecules of the gas, two molecular weights (88) , or two
moles (88 grams) . A formula has an additional meaning which is
important in many chemical calculations.
MOLECULES 283
In dealing with a gas or vapor, it is often necessary to know the
volume that a quantity of it occupies. The unit of measurement of
gas volumes is the volume occupied by one mole (abbreviated M) .
This is called the gram-molecular volume (V) . Study has shown that
the gram-molecular volume is the same for all gases. Under standard
conditions of temperature and pressure, one M of any gas or vapor oc-
cupies 22.4 liters. This may be demonstrated by the experimental
process of weighing a given volume of any gas. For example, one
liter of hydrogen weighs 0.08987 gram. Therefore 22.4 liters weigh
2.016 grams, which is the gram-molecular weight (mole) of hydrogen.
Since one mole of any gas occupies 22.4 liters, we may use the
formula of the gas to represent its gram-molecular volume. For exam-
ple, CO2, which represents one mole of carbon dioxide, also repre-
sents one gram-molecular volume (V) of carbon dioxide; NH3 rep-
resents one M of ammonia and also one V of ammonia. In each case,
V = 22.4 liters. If a coefficient appears in front of the formula of a
gas, it represents the number of gram-molecular volumes. Thus 2CO2
stands for 2V (44.8 liters) ; 2NH3 also stands for 2V. How many liters
of gas would be represented by 4CO,? by
TYPES OF PROBLEMS
TYPE 3B: WEIGHT-VOLUME AND VOLUME-WEIGHT PROBLEMS
In these problems, the weight of one substance is given and
the volume of another is to be found. Or the volume of one
is given, and the weight of another is to be found. The pro-
cedure is the same in both cases. Standard conditions of tem-
perature and pressure (S.T.P.) are assumed.
EXAMPLE: How many liters of nitric oxide can be prepared
by action of sufficient dilute nitric acid on 127.2 g. of copper?
Procedure.
1. Write the balanced equation.
3Cu + 8HNO3 -» 3Cu(NO3)2 4- 4H2O -f 2NO
2. Write the given weight over its formula and x 1. over the
formula whose volume is to be found. Cross out all other
formulas.
127.2 g. x 1.
3Cu +-SHN03 ->4€«fN0i); +-4HsO-+ 2NO
284 NEW WORLD OF CHEMISTRY
3. Under the formula whose weight is given, write its molec-
ular weight. Under the formula whose volume is to be found,
write its gram-molecular volume (V) , not its molecular weight.
127.2 g. x 1.
3Cu -» 2NO
3(63.6) - 190.8 2V - 2(22.4) = 44.8
4. Write the mathematical equation based on the relation-
ship:
wt. of substance used vol. of substance formed
mol. wt. of substance used V of substance formed
Solve for x. 127.2 _ xl.
190.8 " 44.8
190.8* = 127.2(44.8)
x = 29.9 liters of NO
The same general procedure is followed in finding the
weight of one substance when the volume of another is given
except that x represents the unknown weight rather than the
unknown volume. Use an equation based on the same rela-
tionship for your final solution.
PRACTICE WORK ON PROBLEMS OF TYPE 3B
1. What volume of H2 may be obtained by the electrolysis
of 90 g. of H2O?
2. How many liters of NH3 can be formed by the action of
33 g. of (NHJ 2SO4 on sufficient Ca (OH) 2?
3. How much NaCl is needed to produce 112 1. of HC1 gas?
Nad + H2S04 -> NaHSO4 + HC1
4. What weight of H2O must be decomposed to produce
224 ml. of O2?
5. A manufacturer requires 10,000 1. of N2O. What weight
of NH4NO3 must be decomposed?
TYPE 3C: STRAIGHT VOLUME PROBLEMS
This type of problem involves finding the volume of one gas
or vapor when jhe volume of another is known. As we have
learned, the coefficient before the formula of a gas represents
the number of gram-molecular volumes of the gas. Since we
are dealing only with volumes, weights are disregarded. Only
the volumes as represented by the coefficients are considered.
MOLECULES
285
EXAMPLE: How many liters of carbon dioxide are formed dur-
ing complete combustion of seven liters of benzene, C H ?
' 66
Procedure.
1. Write the balanced equation.
2C6H6 + 15O2 -» 12CO2 + 6H2O
2. Write the given volume above its formula. Write x L
above the formula whose volume is to be found. Cross out all
other formulas. , "
7 liters x I.
3. Write the number of gram-molecular volumes (shown by
the coefficients) under the respective formulas.
7 liters x 1.
2C6H6 -» 12C02
2 12
4. Write out the mathematical equation based on the re-
lationship:
_ vol. of substance used vol. of substance formed
coefficient of substance used coefficient of substance formed
Solve for x. 7 _ x 1.
2 "12"
2* = 84
x - 42 liters of CO2
PRACTICE WORK ON PROBLEMS OF TYPE 3C
1. 50 1. of H2 react completely with C12. What volume of
HC1 gas is formed?
2. What volume of H2 is necessary to unite with 5 1. of O2
without leaving any O2 in excess?
3. What volume of NH8 can be made from 5000 1. of pure
N2?
4. How many liters of O2 will be used during the complete
combustion of 500 ml. of methane, CH4?
GH4 + 202 -> G02 + 2H20
5. What volume of oxygen will convert 50 ml. of NO into
nitrogen dioxide, NO2?
286 NEW WORLD OF CHEMISTRY
YOU WILL ENJOY READING
Jaffe, Bernard. Chemical Calculations. World Book Co.,
Yonkers, N.Y., 1947. Additional problems of the types dis-
cussed in this chapter, together with a more detailed account
of methods of determining atomic weights and molecular
weights.
Jaffe, Bernard. Crucibles: The Story of Chemistry f pp. 157-
174. Simon and Schuster, New York, 1948. "The Spirit of a
Dead Man Leads a Battle" tells the story of Avogadro.
Perrin, Jean B. Atoms. D. Van Nostrand Co., New York,
1923. In this book, Perrin, who won the Nobel prize for his
work on the molecule, describes his verification of Avogadro's
law.
USEFUL IDEAS DEVELOPED
1. A molecule is the smallest part of either an element or a
compound which has the properties of that substance.
2. Gay-Lussac's law states that the relation between the
combining volumes of gases and the volumes of their products,
if gaseous, may be expressed in small whole numbers.
3. Boyle's law states that the volume of a gas varies in-
versely as the pressure exerted upon it if the temperature re-
mains constant.
V1/V2 = P2/P!
4. Charles' law states that the volume of a gas varies directly
as the absolute temperature if the pressure on the gas remains
constant.
V1/V2 = Ti/T,
5. Avogadro's law states that equal volumes of all gases
under the same conditions of temperature and pressure are
composed of the same number of molecules.
6. Avogadro's law is valuable because (1) it shows that the
molecules of certain elementary gases, among them hydrogen,
oxygen, nitrogen, and chlorine, consist of two atoms; (2) it
makes possible the determination of the molecular weights of
gases; (3) it makes possible the determination of the atomic
weights of gaseous elements; and (4) it shows the relation-
ships among several apparently conflicting facts concerning the
actions of gases.
7. Avogadro's hypothesis was verified by the work of other
scientists and today is a chemical law. The actual number of
MOLECULES 287
molecules in the gram-molecular weight of a gas was deter-
mined by experiment. This number, called Avogadro's number,
is the same for all gases. It is 6.02 X 1Q28-
8. The gram-molecular volume of a gas is the volume occu-
pied by its gram-molecular weight. Under standard conditions,
it is 22.4 liters.
USING WHAT YOU HAVE LEARNED
Group A
1. In 1860, what was the condition of chemical usage with
respect to the terms atom and molecule?
2. What is the difference between an atom and a molecule?
3. When and by whom was the term molecule first clearly
defined?
4. What did Cannizzaro do to establish the meaning of
molecule?
5. (a) State Gay-Lussac's law and (b) give two illustrations
ofit' '
r
6. The volume of a gas changes from 10 to 5 1. when the
pressure on it changes from 1 to 2 atm. What law does this
illustrate?
7. The volume of a gas changes from 4 1. to 2 1. when its
temperature changes from 500°A. to 250°A. State the law
illustrated.
8. State Avogadro's law.
9. Upon what facts did Avogadro base his hypothesis?
10. State two ways in which Avogadro's law is valuable.
1 1 . What two scientists verified Avogadro's hypothesis?
12. What is a gram-molecular volume of a gas?
13. Outline the method used in working a weight- volume
problem.
14. Outline the method used in working a straight- volume
problem.
15. If 15 1. of N2 are needed to unite with O2 in forming
NO, what volume of O2 will be used?
t . . .
288 NEW WORLD OF CHEMISTRY
16. Assume that air contains 20 percent O2 by volume. What
volume of air will be needed in forming 100 ml. of O3?
17. What volume of air will be needed for the complete
combustion of 750 ml. of acetylene, C2H2?
18. CO passed over warm Ca (OH) 2 reacts as follows:
CO + Ca(OH)2 -» CaCO3 + H2 |
How does the volume of CO compare with that of the H2?
19. What weight of carbon is in 44.8 1. of CO?
20. What volume of NO2 will be formed by the complete
reaction of 100.5 g. of Hg with concentrated HNO3?
Hg + 4HN03 -> Hg(N03)2 + 2H2O + 2NO2
21. HC1 gas was bubbled through a solution of NaOH. As
a result, 468 g. of NaCl were formed. What volume of the
HC1 used actually combined with the base?
Group B
22. From the experimental fact that 3 vol. of O2 change into
2 vol. of O3 when an electric discharge is passed through moist
O2, prove that the molecule of ozone contains three atoms.
23. Prove nitrogen molecule contains at least two atoms.
24. Which is more economical to use in the preparation of
NH3, (NH4) 2SO? at $8.75 per 100 Ib. or NH^Cl at 12^ per lb.?
25. (a) Explain the operation of the atomic-hydrogen torch,
(b) Account for the extreme heat obtained.
26. What is meant by absolute temperature?
27. Prove by calculation that the ounce-molecular-volume
of any gas equals 22.4 cu. ft.
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. In England, Robert Boyle is considered the father of
modern chemistry. In France, Lavoisier is called the father of
modern chemistry. Can you suggest reasons for this difference
of opinion among French and English scientists? Is it justified?
Does it illustrate scientific open-mindedness?
2. Suggest possible reasons for the neglect of Avogadro's
hypothesis from 1811 to 1860. Can you cite other scientific
work which remained unrecognized for a long time?
3. Construct a cardboard or wooden box to represent the
gram-molecular volume of any gas at S.T.P. Exhibit it to your
class alongside of a quart bottle or carton of milk.
20
SULFUR
AND
HYDROGEN SULFIDE
. . . Sulfur has been taken intermit-
tently from Popocatepetl since the
time of the ancient Aztecs, who used
it for medicinal purposes. Two of
Cortez's soldiers, in the sixteenth
century, climbed to the crater and
obtained sulfur for the purpose of
manufacturing gunpowder. Science
News Letter
An American pharmacist creates a new industry. The discovery
of petroleum in Pennsylvania in 1859 led at once to a wide search
for other stores of oil. Only six years later, oil prospectors stumbled
upon huge deposits of sulfur, a yellow, brittle, lustrous solid known
since ancient times. These deposits were about 500 feet below the
surface in Louisiana not far from the Gulf near the Texas line. Here
ages ago a vast geyser had spouted, leaving the sulfur within and
about its crater. The sulfur was covered with strata of clay, limestone,
and, worst of all, gas and quicksand. It was impossible to sink shafts
to reach the deposits in order to dig the sulfur out as coal is mined.
Many companies were formed to exploit these deposits, but because
of the many difficulties, they all failed.
In 1891 Herman Frasch (frash) heard about this sulfur. He had
come from Germany 25 years before. Leaving high school at 16, he
had been apprenticed to a druggist and then had left for the United
States, where he opened his own drugstore in Philadelphia. Chem-
istry fascinated him, and in the back of his drugstore he carried on
many researches on petroleum products. Later he sold his store and
devoted all his time to chemical engineering. He tackled the problem
of extracting sulfur from the Louisiana deposits.
289
compressed
u alr
sulfur
superheated
water
water
melted sulfur
' '•' :• /' :' V"'.\ : '••'.'•" anhydrite • "• / ,'v
Fig. 64. The Frasch process for extracting sulfur from deposits deep in the earth.
When Frasch described to some of his friends the entirely new
process that he had devised, they thought it impossible. One promi-
nent person challenged him in no uncertain terms. He said that he
would swallow every ounce of sulfur Frasch extracted by his queer
process of pumping a solid out of the earth. But Frasch kept on im-
proving his method and ultimately he succeeded in founding a new
American industry based on his process.
The Frasch process of extracting sulfur. Frasch' s plan was to sink
a well by means of an oil-drilling rig, and lower three concentric
pipes (inside a casing) down to the sulfur. Through the outermost
6-inch pipe, superheated water was to be pumped to melt the sulfur.
Through the innermost one-inch pipe, compressed air was to be
forced down to the sulfur deposit to make the melted sulfur frothy
and light. The result was to be a flood of molten, foamy sulfur gush-
ing under pressure from the three-inch pipe between the other two.
Frasch was visibly nervous when he gave the order to start the
first pump. More and more slowly went the engine with its increas-
ing load until the man at the throttle sang out at the top of his voice,
"She's a-pumping." A liquid appeared at the mouth of the three-inch
pipe. Frasch wiped the liquid off a polished piston rod, and gazed
upon the first crystals of sulfur extracted by his ingenious process.
Then came a steady stream of golden liquid, which, in 15 minutes,
filled every one of the 40 barrels available. Still the molten sulfur
kept pouring. Embankments were quickly thrown up and lined with
boards to hold the sulfur as it solidified.
This is the method of mining sulfur in use today in Texas and
Louisiana. These deposits now supply much of the world's needs. In
this way, mountains of sulfur, 99^ percent pure, are extracted and
stand ready to be dynamited into pieces small enough to be loaded
and shipped to many parts of the world.
290
SULFUR AND HYDROGEN SULFIDE 291
Sulfur is obtained in Sicily by a different process. Before the time
of Frasch, a group of English financiers had been marketing the rich
sulfur deposits of the volcanic region of Sicily. Here the sulfur is
found mixed with clay and limestone, from which it is separated by
melting this ore, allowing the free sulfur to flow away from the im-
purities. The sulfur is then boiled and is changed to a powder, called
(loiuers of sulfur, by chilling the vapors.
The English monopoly had been supplying more than 90 percent
of the world's sulfur. In 1904, when American output reached the
point where a single well could supply 400 tons of sulfur daily, the
English company went out of business. To prevent the unemploy-
ment of hundreds of workers, the Italian government took control.
Frasch aided gladly in stabili/ing the Sicilian sulfur industry.
Crude free sulfur accounts for about only 40 percent of the world's
sulfur production. Iron, zinc, and copper sulficles, natural gas, and
industrial gases supply the rest. The United States today leads the
world in the production of sulfur. Many countries, including Italy
and Mexico, are stimulating their own production of sulfur by sub-
sidies, tariff laws, bounties, and price guarantees.
Physical properties of sulfur. Sulfur is a pale yellow, soft, brittle
solid (plastic, or pliable, in the case of amorphous sulfur) with a
very faint odor and no marked taste. It is practically insoluble in
water, more soluble in carbon tetrachloride, CC14, and very soluble
in carbon disulfide, CS...
Sulfur is a poor conductor of heat. It melts at about 114°C.,
forming a pale yellow liquid, which on further heating darkens and
thickens, becoming- almost black at 235 °C. At a still higher tempera-
ture, it becomes thinner again, and finally changes to a yellow vapor
at 445°C. Sulfur is mentioned in the Bible as brimstone.
Texas Gulf Sulfur Cohipany
Liquid sulfur flows into the
storage vat in which it will
solidify. The pipeline leads
to the sulfur well.
292
NEW WORLD OF CHEMISTRY
Allotropic forms of sulfur. Oxygen, as you have learned, occurs in
two allotropic forms, O2 and O3, but sulfur occurs in two common
crystalline forms (rhombic and prismatic) , and one noncrystalline
or amorphous form. Each has different properties caused by differ-
ences in the arrangement of their atoms.
1) Rhombic sulfur. Sulfur is found in nature in the form of
rhombic crystals, the most stable form under normal conditions.
Their molecules consist of rings of eight atoms of sulfur (Ss) , and
their density is about two. Crystals of this form of sulfur may be
prepared by dissolving sulfur in carbon disulftde and allowing the
solvent to vaporize slowly. The residue consists of perfect crystals
having the shape of two pyramids with their bases joined. Roll sulfur,
made by cooling sulfur in cylindrical molds, is almost entirely
rhombic.
2) Prismatic sulfur. When sulhir is heated until it just melts and
is then allowed to cool slowly, it forms long needle-shaped crystals
whose density is somewhat less than that of rhombic sulfur.
3) Amorphous sulfur. When boiling sulfur is suddenly cooled by
pouring it into cold water, the rings are broken and are replaced by
long chains of sulfur atoms which become entangled and can be
stretched like rubber. It is amber in color and, unlike the other
two forms, is insoluble in carbon disulfide. Amorphous sulfur
changes gradually into rhombic at room temperature. For simplicity,
all forms are designated S. Flowers of sulfur is a powdered mixture
composed of rhombic and some of the plastic. Between 96 °C. and
1 14°C. the most stable form of sulfur is the prismatic.
Chemical properties of sulfur. The atomic weight of sulfur is
32. Its atomic number is 16, and it has, therefore, six electrons in its
third, or outermost, ring. This makes it a borrower of electrons. It
is a nonmetal, fairly active, and has, under ordinary conditions, a
valence of minus two. Therefore, sulfur closely resembles oxygen in
its chemical properties.
Fig. 65. Allotropic form, of sulfur. melfed Sulfur"
rhombic
amorphous
sulfur
Exterior of a sulfur storage vat.
The solidified sulfur will be
blasted into small pieces for
shipment in railroad cars.
Texas Gulf Sulfur Company
Like oxygen, sulfur unites with most metals. The compounds of
metals and sulfur are called sulfides. For example, when sulfur is
heated with iron powder, iron sulfide is formed.
- - • Fe-fS-^FeS
Hot copper burns in sulfur vapor, forming cuprous sulfide.
. /- V^ • • 2Cu+S-+Cu2S
When sulfur is mixed with zinc dust and ignited, the chemical
reaction is so vigorous that a great deal of light and heat are liber-
ated and dense clouds of zinc sulfide, which settle out as a powder,
are formed.
Zn + S-^ ZnS
Although sulfur is a nonmetal, it is less nonmetallic than oxygen.
It therefore combines with oxygen, exhibiting a positive valence of
either four or six. Sulfur burns in air with a pale blue flame, form-
ing sulfur dioxide, SO.,. A small amount of sulfur trioxide, SO3, is
formed later by further oxidation of the SO2.
2SO2 + O2
> SO2 (sulfur dioxide)
> 2SO3 (sulfur trioxide)
The electron structure of sulfur and oxygen are shown below.
Fig. 66. The structure of the oxygen atom (left) and the sulfur atom.
Oxygen
Sulfur
293
294 NEW WORLD OF CHEMISTRY
Sulfur also shows more metallic properties than chlorine by unit-
ing with the latter, forming sulfur dichloride, SCI.,, a brownish red
liquid with a pungent odor used in chlorinating, and also sulfur
monochloride, a heavy, amber-colored unstable liquid with an irri-
tating odor, used in vulcanizing rubber.
2S + C12 -> S2C12 (sulfur monochloride)
When sulfur vapor is passed over carbon heated in an electric fur-
nace, the two elements combine, forming carbon disulfide, CS2.
C + 2S -» CS2
Carbon disulfide is a heavy, colorless liquid with a pleasant odor.
Generally, however, it has a disagreeable odor caused by impurities.
It is very combustible. Its chief use is as a solvent for sulfur, gums,
rubber, fats, and waxes. It has been used also as a poison in exter-
minating ants and other insects, rats, and prairie dogs.
Rubber tires depend on sulfur. The normal annual consumption
of sulfur in the United States is more than three million tons. It is
one of the fundamental industrial elements. By far the greatest
quantities of sulfur are used in the manufacture of sulfur ic acid,
one of the most widely used of the heavy chemicals. The production
and uses of sulfuric acid are discussed in Chapter 21.
Great quantities of sulfur are used in vulcanizing rubber. By this
process, soft, sticky, perishable, natural rubber is changed to a harder,
less plastic, more durable rubber. Vulcani/ed rubber is used chiefly
in making automobile tires but also in making thousands of other
rubber articles.
Tfie Firestone Tire and Rubber Company
In a modern tire factory, a worlc-
man removes a finished tire
•om a steam-heated vulcani-
zing mold.
SULFUR AND HYDROGEN SULFIDE 295
When the American, Charles Goodyear, in 1839 while working in
his kitchen in Woburn, Massachusetts, accidentally dropped a piece
of rubber mixed with sulfur on a hot stove, he discovered the proc-
ess of vulcanization (derived from Vulcan, the Roman god of fire) .
A new and highly important industry was made possible. To shorten
the time required for vulcanizing rubber, a catalyst, or accelerator,
such as zinc oxide, is added to the mixture of rubber and sulfur.
In making an article of rubber, the washed raw rubber is thor-
oughly mixed with various chemicals that determine the properties
of the finished product. Among these substances are sulfur, carbon
black, lead oxide (PbO) , zinc oxide, and carbonates of magnesium
and calcium. Different combinations of these and other substances
in varying quantities may be used in accordance with the properties
desired in the finished product. The rubber is then rolled into sheets
of the desired thickness or placed in molds of the desired shape.
While in the mold, the rubber is heated and vulcanization occurs.
Sulfur is used in controlling fungi and insects. A very effective
liquid for destroying fungus growths and fruit tree, shrub, and vine
pests is a deep orange-red lime-sulfur spray made by boiling sulfur
with calcium hydroxide. The spray is both a fungicide, a substance
that kills fungi and molds, and an insecticide, a substance that kills
insects. Dusting with very finely powdered sulfur is effective against
rose diseases, mildew, and black spot. However, ordinary flowers of
sulfur is not fine enough to be of much value as a dusting agent, and
even some finely ground commercial dusting powders are too coarse.
Colloidal sulfur and wettable sulfur, the first a very highly dis-
persed sulfur in water and the second sulfur so treated that it dis-
perses on contact with water, are both used in making mild sprays
that are particularly useful for the summer spraying of roses and for
the control of mildew and true rust diseases of other plants.
Sulfur is used in medicine. Ointments containing sulfur have
been used since antiquity to control skin diseases caused by fungi.
A common ointment of this kind is made by mixing three parts by
weight of sulfur, 15 parts of white petrolatum, or Vaseline, one part
of lanolin, or wool fat, and one part of yellow wax. Such an oint-
ment is effective in killing the very small mites that cause scabies,
or itch. Sulfur ointments may be of value in the control of infectious
dandruff.
Other uses of sulfur. Much sulfur goes into the manufacture of
calcium bisulfite, Ca (HSO3) 2, which is used in making wood pulp
for the manufacture of paper. Sulfur is used also in making syn-
thetic resins, sulfur colors, and gunpowder. Sulfur-lined steel pipes
296
NEW WORLD OF CHEMISTRY
are used to transport liquids that are very corrosive to the materials
of which pipes are normally made. Sulfur cements are used to join
bricks in floors and walls that are continually subjected to the corro-
sive effects of acids or alkalies. Great quantities of sulfur are used
in the making of matches.
How matches are made. Early matches were dangerous, hard to
use, and hard to carry. Phosphorus matches often caused painful
burns and were harmful to the health of the workers who made them.
The first friction match was invented in 1827 by an English phar-
macist, John Walker. It contained potassium chlorate (KC1O3) , an
oxidi/ing agent, and antimony sulfide (Sb,Ss) , a compound with a
low kindling temperature. The locofoco match was an American
adaptation of Walker's friction match. During the presidential cam-
paign of 1840, the Democrats were called locofocos because at one of
their meetings, they used matchlight when the Whig landlord turned
off the gas.
In 1831 white phosphorus was used for the first time in making
matches. It was a more efficient fire-producer than antimony sulfide,
but played havoc with the health of the match-factory workers, who
finally demanded protection. This led in 1906 to an international
convention which prohibited the further use of white phosphorus
in making matches because of its poisonous nature.
The strike-anywhere match in use today also contains both an oxi-
dizing agent and a compound with a low kindling temperature. The
head of the paraffin-dipped matchstick contains potassium chlorate
and phosphorus sulfide (P4S3) , a dark solid that is concentrated on
the tip of the match. Glue is used to bind the chemicals, and ground
glass or other abrasive is sometimes added as a filler.
The tip of a safety match is composed of easily combustible anti-
mony sulfide (Sb.,S3) , and potassium chlorate, which provides addi-
tional oxygen. The side of the box contains red phosphorus, which
H
H
hydrogen sulfide S
Fig. 67. Laboratory preparation of
hydrogen sulflde (left). Structure
of the hydrogen sulflde molecule
(right).
is nonpoisonous. The material on the tip of the match will not ig-
nite easily unless it is rubbed on the treated side of the box.
Disastrous fires are frequently caused by matches that have been
blown out but still retain glowing tips. To prevent this dangerous
afterglow, matches are dipped in some such solution as sodium sul-
late or ammonium phosphate during their manufacture. When
blown out, they do not leave glowing tips.
Waterproof matches are treated with a transparent coating with
a high-kindling temperature. They will light even after being soaked
in water.
The manufacturer of matches uses a continuous-process machine,
which takes pine wood, cuts it, dips the sticks in paraffin, coats them
with the chemicals needed, dries them, and finally packs them for
shipment. More than a million matches can be made by one of
these machines each hour.
How sulfur and hydrogen unite. Hydrogen has only one electron.
It therefore needs one electron to complete its only ring; but since
it shows a strong tendency to lend its one electron, it is said to pos-
sess metallic properties. Sulfur, requiring two electrons to complete
its outer ring, will combine with two atoms of hydrogen. This union
forms the compound, hydrogen sulfide, H,S, as shown in Fig. 67.
In this compound, sulfur exhibits definitely nonmetallic properties.
Although it is possible to prepare hydrogen sulfide by direct union
of sulfur and hydrogen, it is prepared most easily by other methods.
298 NEW WORLD OF CHEMISTRY
How hydrogen sulfide is prepared. Hydrogen sulfide is prepared
both commercially and in the laboratory by the general method of
preparing an acid, that is, by the action of sulfuric acid on a salt (sul-
fide) . The sulfide most generally used is ferrous sulfide, FeS, a black
iron ore. When sulfuric acid is poured on ferrous sulfide in a test
tube (see Fig. 67) , hydrogen sulfide gas is given off immediately
without the addition of any external heat.
FeS -I- H2SO4 -> FeSO4 + H2S |
Physical properties of hydrogen sulfide gas. Hydrogen sulfide is
colorless, slightly heavier than air, and fairly soluble in water (one
volume of water dissolves three volumes of hydrogen sulfide) . Natu-
ral sulfur waters contain hydrogen sulfide in solution and, upon be-
ing decomposed, leave a deposit of free sulfur. Easily liquefied,
hydrogen sulfide is sold in cylinders for laboratory use. Its most char-
acteristic physical property is its odor, the odor of rotten eggs. In
fact, it is hydrogen sulfide that gives such eggs their odor. It is caused
by the decomposition of organic sulfur compounds in the yolks.
Hydrogen sulfide forms naturally in marshes, oil wells, mines and
coal piles, manure pits, and sewers. In persons who have endured
mild exposure to its effects, it produces inflamed thrpat, headache,
a heavy feeling in the stomach, and dizziness. When breathed in
large quantities it causes death. Both natural gas and coal gas con-
tain H2S which is removed before they are used as household fuels.
Chemical properties of hydrogen sulfide gas. When burned in
sufficient air, hydrogen sulfide gives off a pale blue flame, and water
vapor and sulfur dioxide are formed. This sulfur dioxide gives burn-
ing hydrogen sulfide its irritating odor. The equation for the com-
plete combustion of hydrogen sulfide is:
2H2S + 3O2 -» 2H2O + 2S02 T
When hydrogen sulfide is burned in a small amount of air (incom-
plete combustion) , water is formed as before, but free sulfur is pro-
duced instead of sulfur dioxide. This sulfur separates out as a yellow
powder. The fact that it does so probably accounts for the presence
of sulfur around volcanoes, which emit hydrogen sulfide gas.
2H2S + O2 -» 2H2O + 2S T
Because of the tendency of hydrogen sulfide to unite with oxygen,
it is a fairly good reducing agent. When hydrogen sulfide acts as a
reducing agent, the sulfur lends electrons and is oxidized.
SULFUR AND HYDROGEN SULFIDE
299
In a water solution, hydrogen sulfide dissociates to some extent.
Such a solution acts as a weak, unstable acid, sometimes known as
hydrosulfuric acid., H2S. On continued boiling, this acid liberates hy-
drogen sulfide, leaving pure water.
Hydrogen sulfide reacts with certain metals and also with the salts
of certain metals, forming sulfides. The tarnishing of silverware is
caused by the formation of a brownish-black sulfide of silver, Ag2S.
The blackening of lead paints is caused by the formation of black
lead sulfide. Lithopone, a white paint base now used widely, consists
of a mixture of barium sulfate (RaSO4) , and zinc sulfide (ZnS) .
Lithopoiie paints do not lose their color by the action of H,S. Why?
Titanium dioxide, TiO2, also is used as a white paint base. It is not
blackened by sulfur compounds and has great covering power.
Many important sulfides are found in the earth. The salts of
hydrosulfuric acid form an important class of compounds called
sulfides. Many of them occur in nature and constitute important ores
such as iron pyrites., FeS.,; galena, PbS; zinc blende, ZnS; cinnabar,
HgS; and CuS. Certain colored sulfides are used as mineral pig-
ments in the coloring of paints. Cadmium sulfide, CdS, for example,
is a yellow pigment, and zinc sulfide, ZnS, is a white pigment used
in paints (see illustration following page 382).
The chief use of hydrogen sulfide. In the analysis of ores and in
the separation of groups of certain metals from other groups of met-
als, hydrogen sulfide gas is indispensable. For this reason, a hydro-
gen sulfide generator is always present in a laboratory for analytical
chemistry. The sulfides of certain metals such as sodium and calcium
are soluble in water, while those of other metals such as lead and zinc
are insoluble.
By passing hydrogen sulfide into a solution of the soluble salts of
such metals, the sulfides of certain metals precipitate out and may
be separated by filtration. Thus, certain metals present in an ore or
Standard Oil Com puny (N .J .)
Hydrogen sulfide fumes ris-
ing from crude oil may
prove fatal to workmen.
Hence, protective masks
are worn by anyone work-
ing near the oil storage
tanks.
300 NEW WORLD OF CHEMISTRY
alloy may be separated with the aid of H2S. Furthermore, since the
colors of the sulfides of many metals differ, chemists can use these
differences in color in detecting the presence of these metals. For
example, zinc sulfide is white, arsenic sulfide is yellow, antimony
sulfide is orange, and copper sulfide is brownish-black.
Zn(NO3)2 + H2S -> ZnS I + 2HNO3
CuSO4 + H2S -> CuS I + H2SO4
This difference in color is only one of the very many ingenious
methods used by analytical chemists in detecting and isolating ele-
ments present in complex mixtures and compounds.
The test for the sulfide ion. If, on the addition of sulfuric acid to a
compound, hydrogen sulfide gas is liberated, the substance tested
is a sulfide. The hydrogen sulfide liberated is easily detected either
by its odor or by its ability to turn a silver coin brownish-black, as a
result of the formation of silver sulfide.
YOU WILL ENJOY READING
Fabre, Jean H. The Wonder Book of Chemistry, pp. 345-359.
Albert &T Charles Boni, New York, 1922. Discusses sulfur and
includes some simple experiments with this element.
Waggaman, W. H., and Barr, J. A. "Sulfur for Survival."
Chemistry, October, 1951, pp. 1-10. An illustrated article on
sources, extraction, and properties of sulfur.
USING WHAT YOU HAVE LEARNED
Group A
1. When and by whom was the first successful method of
obtaining S from the Louisiana deposits invented?
2. What was the great difficulty that had to be overcome
before S could be extracted from the Louisiana deposits?
3. How did the Frasch process affect the Sicilian sulfur
industry?
4. (a) Make a labeled diagram of the Frasch process and
(b) explain it.
5. What are two ways in which the Louisiana sulfur de-
posits differ from the Sicilian sulfur deposits?
6. Name the allotropic forms of S and tell how each may
be prepared in the laboratory.
SULFUR AND HYDROGEN SULFIDE 301
7. What two elements other than S occur in allotropic
forms?
8. (a) State the chief differences and resemblances of the
three allotropic forms of S. (b) How could you prove that all
three forms of S are the same element?
9. A piece of plastic S is left overnight in the laboratory.
The next morning yellow brittle S is found. Explain.
10. (a) How are the two kinds of commercial sulfur, roll
sulfur and flowers of sulfur, prepared? (b) What allotropic
forms are in each?
1 1 . Make a diagram of an atom of S and use it to find the
valence of S, to describe its chemical activity, and to explain
why it is a nonmetal.
12. In what three ways does S resemble O2?
13. Write balanced equations for the: (a) union of sulfur
and zinc; (b) complete combustion of hydrogen sulfide; (c) in-
complete combustion of hydrogen sulfide; (d) burning of sul-
fur; (e) union of sulfur and chlorine; (f) complete combustion
of carbon disulfide.
14. What are four chief uses of S? List in order of im-
portance.
t . . .
15. (a) What is meant by the vulcanizing of rubber?
(b) What properties does vulcanizing impart to rubber?
(c) What American discovered this process?
16. (a) Of what chemicals is the tip of a strike-anywhere
match composed? (b) How does a safety match differ from an
ordinary match?
17. Hydrogen sulfide present in natural gas is removed be-
fore the natural gas is sent into the pipelines to be used as fuel.
Explain.
18. By a labeled diagram describe the laboratory prepara-
tion of H2S.
19. What weight of FeS would be needed to prepare 204 Ib.
of H2S?
20. What volume of gaseous H2S could be prepared from
20 g. of 90 percent pure FeS?
21. A compound of hydrogen and sulfur has a mol. wt. of 34.
The percentage of S in the compound is 94.1 percent. Find its
formula.
302 NEW WORLD OF CHEMISTRY
22. What volume of H2S is required to precipitate all the
CuS from a solution containing 80 g. of CuSO4?
23. What volume of SO2 will be formed by the complete
combustion of 896 ml. of gaseous H2S?
24. Write ionic equations for the following and tell whether
each reaction goes to completion (refer to Table 10, page 243) :
a) FeS + H2S04 -» FeSO4 + H2S
b) Pb(N03)2 + H2S -» PbS + 2HNO3
c) ZnSO4 + (NH4)2S -> ZnS + (NH4)2SO4
d) Na2SO4 + H2S -» Na2S + H2SO4
25. Compare the physical properties of H2S and N2O.
26. Describe what happens when gaseous H2S is bubbled
through water.
27. Why is H2S said to be (a) weak? (b) unstable?
28. Explain the tarnishing of silverware.
29. Why is H2S solution kept in amber-colored bottles?
30. Using two equations, show the chief use of H2S.
31. Describe a test for a sulfide.
32. Copy and complete. Do not write in this book. S occurs
in our bodies because it is one of the elements found in ....
The two chief sulfur-producing states in America are Louisi-
ana and .... A lime-sulfur preparation is used as .... A sulfur
compound used in exterminating rats is .... A sulfide that is
soluble in water is ....
33. Under a rubber band used to keep pieces of silver
cutlery together, a black mark is found. Explain.
Group B
34. How would you tell fool's gold, FeS2, from genuine gold?
35. In the Frasch process, why is S forced out of the three-
inch pipe rather than out of the six-inch pipe?
36. The action of H2S on a solution of ZnSO4 is reversible.
How would you force this reaction to completion?
37. NH3 leaks may be detected by burning sulfur candles.
Explain.
38. (a) Why can we not use HNO8 in preparing H2S from
FeS? (b) What is formed if H2S is passed into HNO3?
SULFUR AND HYDROGEN SULFIDE 303
39. Why do we write CS2 and not S2C?
40. A mixture of sulfur and molasses was given to children
as a "spring tonic." What do you think of this practice?
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Discovery of the vulcanization of rubber is generally con-
sidered an accident. Had the accident not happened to Good-
year, would the thousands of rubber articles in use today never
have come? Explain your answer fully in terms of scientific
advance and the needs of society.
2. If you live in a rural section, talk with as many farmers
as possible to learn what methods they use to check the losses
incurred by insect pests and fungus growths. Report your
findings to your class.
3. Filter paper dipped in a solution of Pb (C2H3O2) 2 turns
black in the presence of H2S. Hang small strips of it in the
basement while the furnace is burning, over the kitchen stove
or gas range while dinner is being cooked, and in the living
room. Which of the strips turns blackest first? What do you
conclude?
21
SULFURIC ACID:
THE FUNDAMENTAL ACID
. . . We may fairly judge of the
commercial prosperity of a country
from the amount of sulfuric acid it
consumes. Reflecting upon the im-
portant influence which the price of
sulfur exercises . . . we can under-
stand why the English Government
should have resolved to war with
Naples, in order to abolish the siu-
fur monopoly, which the latter
power attempted recently to estab-
lish. Justus yon Liebig,
Why sulfuric acid is the fundamental chemical today. One of the
outstanding differences between our society and the society which
preceded it is the tremendous role that machines play. Automobiles,
locomotives, steel ships, and airplanes are comparatively new. In the
manufacture, operation, and upkeep of these and thousands of other
machines, sulfuric acid is directly or indirectly a prime necessity. If
the ability to make this acid were lost suddenly, industry would be
paralyzed. So fundamental is it to our industrial life that its con-
sumption is a fair index of industrial conditions.
How sulfur dioxide is prepared for use in industry. About 90 per-
cent of the sulfur consumed industrially is first burned in air to
produce sulfur dioxide. Heating sulfide ores is another commercial
method of preparing SO2. When an ore, such as iron pyrites, FeS2, is
heated in air, or roasted, one of the products formed is sulfur dioxide.
4FeS2 + 11O2 -» 2Fe2O3 + 8SO2 T
Sulfur dioxide occurs in small quantities in the vapors rising from
active volcanoes and certain sulfur Springs, and in the gases formed
during the burning of coal.
304
Na2SO3 + H2SO4
Fig. 68. Laboratory prepa-
ration of sulfur dioxide.
Why is the excess gas
passed through water?
-*— H2Q
How sulfur dioxide is prepared in the laboratory. When sodium
sulfite, Na2SO3, or any other sulfite is treated with an acid, sulfur
dioxide is liberated. The gas is collected by the displacement of air,
as shown in Fig. 68.
Na2SO3 + H2SO4 -» Na2SO4 + H2O + SO2 1
Physical properties of sulfur dioxide. Sulfur dioxide is colorless,
has a suffocating odor, is more than twice as heavy as air, and is very
soluble in water (one volume of water at room temperature dissolves
40 volumes of the gas) . Under a pressure of only three atmospheres
(approximately 45 pounds per square inch, or 45 psi) it changes to
a colorless liquid, which can be transported in steel cylinders or
shipped like gasoline in tank cars.
Chemical properties of sulfur dioxide. Sulfur dioxide does not
burn in air, but it may be made to combine with another atom of
oxygen, forming sulfur trioxide, SO3, by passing a mixture of sulfur
dioxide and air over a catalyst, such as platinum or vanadium pent-
oxide.
2SO2 + O2 -* 2SO3
This ability of sulfur dioxide to combine with more oxygen makes
sulfur dioxide a reducing agent. Sulfur dioxide may also act as an
oxidizing agent. When, for example, it is passed through a solution
of hydrosulfuric acid, H2S, it precipitates sulfur. The sulfur (S++++)
in the SO2 takes electrons from the sulfur (S — ) of H2S and changes
to free sulfur, S°.
S++++O2— + 2H2+S— -» 2H2+O— + 3S° j
Hydrosulfuric acid acts here as the reducing agent. These two actions
of sulfur dioxide illustrate the fact that a substance may be either an
oxidizing agent or a reducing agent, depending upon the substance
to be reduced or oxidized.
305
306
NEW WORLD OF CHEMISTRY
freezer
suction
refrigerant
check valve-
condenser
shut-off
valve
•H^lpfc^ expansion valve
freezing tray
Fig. 69. An electric refrigerator.
The refrigerant may be sulfur
dioxide, methyl chloride, Freon,
etc.
condenser
electric motor compressor
When sulfur dioxide dissolves in water, it combines with some of
the water, forming sulfurous acid. Sulfur dioxide, then, is an acid
anhydride.
SO2 + H2O — > H2SO3 (sulfurous acid)
Sulfur dioxide has many uses. The most common use of sulfur
dioxide is in the manufacture of sulfuric acid. It is used also to some
extent instead of ammonia in mechanical refrigeration. Although
liquid sulfur dioxide during evaporation absorbs only one-third as
much heat as ammonia does, it liquefies under a lower pressure and
hence is used in some household refrigerators.
Homer in his Odyssey cited the use of burning sulfur in fumiga-
tion. Today rooms, houses, or warehouses are sometimes fumigated
by burning sulfur candles or by liberating sulfur dioxide gas from
cylinders. However, as a fumigant sulfur dioxide has been almost
wholly supplanted by more effective substances. It is even extremely
doubtful that fumigation by sulfur candles was ever very effective.
How sulfurous acid, H,SO3, is prepared in the laboratory. In
the laboratory preparation of sulfur dioxide, SO2 is not the product
first formed. Sulfurous acid is the first product of the reaction, but
it>J£ impossible to stop the reaction at this point.
Na2SO3 + H2SO4 -» Na2SO4 + H2SO3
Sulfurous acid is prepared by passing sulfur dioxide gas through
cold water.
H2O + SO2 -» H2SO3
SULFURIC ACID 307
Properties of sulfurous acid. Sulfurous acid is a colorless solution
with a suffocating odor (sulfur dioxide) . It is unstable and decom-
poses readily into water and sulfur dioxide. By boiling sulfurous
acicl, all the sulfur dioxide may be driven off. The reaction is re-
versible, depending upon the temperature.
H20 + SO. t± H2SO3
Sulfurous acid is a very weak acid because it dissociates only
slightly, forming few hydrogen ions. When neutralized with a base,
it forms salts known as sulfites.
H2SO, + 2NaOH -> Na2SO3 (sodium sulfite) + 2H2O
On exposure to air, sulfurous acid may be oxidized to sulfuric acid
to some extent.
2H2SO3 + O2 -» 2H2SO4
In a similar way, sulfites may be oxidized to sulfates.
How sulfurous acid, sulfites, and bisulfites are used. The tend-
ency of one molecule of sulfurous acid to unite with one atom of
oxygen, forming sulfuric acid, makes it a valuable reducing agent.
This property of sulfurous acid enables it to be used in the bleach-
ing of wood pulp, straw, silk, feathers, dried fruits, flour, molasses,
and canned corn — all of which would be partially destroyed by
other bleaching agents.
Fruits that are to be dyed are frequently bleached with sulfur
dioxide (as sulfurous acicl) . For example, in making maraschino
cherries, the cherries are first bleached and then dyed any desired
color with edible dye. If sulfurous acid is used either in bleaching
or in preserving a food that is sold in interstate commerce, this fact
must be clearly stated on the label of the package in which it is sold.
Chemical Construction Corporation
A sulfur combustion furnace
for the commercial production
of sulfur dioxide. Melted sulfur
is pumped into the furnace
and burned to form the gas.
308 NEW WORLD OF CHEMISTRY
Substances of vegetable origin are not easily destroyed by sulfurous
acid, and therefore their coloring matter may be reduced to colorless
compounds by this weak acid. However, the bleaching action of
sulfurous acid is not as lasting as that of chlorine or hydrogen per-
oxide, because the coloring matter that has been reduced may, on
exposure to air, become reoxidized. This explains in part why straw
hats, bleached with sulfurous acid, may turn brown.
Calcium bisulfite, Ca (HSO3) 2, containing dissolved SO2, is used in
converting ground wood pulp into sulftte pulp by dissolving out the
gluelike substances found mixed with the cellulose of the wood pulp
(see page 598) . Sulfite pulp is used in making paper.
Sodium sulftte mixed with a solid acid, such as oxalic (white crys-
tals) , makes up the dry powder used in bleaching straw hats. The
action of this mixture of acid and sulfite illustrates the part that ions
play in chemical reactions. While the dry solid acid and sulfite are
in contact, no action occurs. As soon as water is added, however, the
substances go into solution, dissociate, and the ions formed react,
producing sulfurous acid, which bleaches. Na2SO3 is also used as an
antichlor (a substance used to remove any excess of chlorine after
bleaching) and as an oxygen-removal agent in treating boiler-feed
water.
How to test for a sulfite. The characteristic od6r of the sulfur
dioxide liberated when an acid is added to a sulfite constitutes a
test for any sulfite.
CaSO3 + 2HC1 -> CaCl2 + H2O + SO2 T
Preparation and properties of sulfur trioxide. When sulfur di-
oxide is mixed with air and the mixture is passed over a catalyst
such as heated platinum, sulfur trioxide is formed.
2SO2 + O2 -> 2SO8
This constitutes the key reaction in the most important commercial
method of manufacturing sulfuric acid, the contact process, so named
because contact of the gases with this platinum catalyst plays so im-
portant a part.
Sulfur trioxide is a white solid, but it melts at about room temper-
ature and boils at 46 °C. The liquid fumes in moist air and reacts
vigorously with water making a hissing sound, liberating much heat
and forming sulfuric acid. Thus, sulfur trioxide is the anhydride of
sulfuric acid.
SO3 + H2O -» H2SO4
•..
,"" for '"«'
!::;Mef
Monsanto Chemical Company and Davison Chemical Corporation
A contact process installation. Trace the process in this photograph and that on
page 311.
How sulfuric acid is made by the contact process. The manu-
facture of sulfuric acid by the contact process may be conveniently
divided into three stages: (1) burning sulfur to sulfur dioxide in
a combustion furnace, (2) changing this sulfur dioxide to sulfur tri-
oxide by passing sulfur dioxide and air over a catalyst in a converter,
(3) changing the sulfur trioxide to sulfuric acid by passing the sul-
fur trioxide into absorption towers through which sulfuric acid is
circulated. The chemical reactions that take place are:
1) S + O2 -> SO2
2) 2S02 4- 02 -> 2S03
3) SO3 4- H2O [+ H?SO4! -» i2|H2SO4
1) The sulfur combustion furnace. In the combustion furnace,
sulfur is brought into intimate contact with thoroughly dried air.
A blower supplies this air at sufficient pressure to force it through
the furnace and to force the sulfur dioxide from the furnace through
the rest of the system. The heat produced by the combustion of the
sulfur is recovered by means of a waste-heat boiler. Part of this heat
is used in some plants to melt the sulfur that is sprayed into the
furnace.
309
310 NEW WORLD OF CHEMISTRY
2) The converter. The converter is simply a large chamber con-
taining many perforated shelves covered with a catalyst in whose
presence the union of warm sulfur dioxide and oxygen takes place,
forming sulfur trioxide. Noteworthy developments in the contact
process have been the introduction of vanadium pentoxide and
platinized silica gel catalysts as substitutes for the more expensive
platinum mass. The vanadium catalyst has the further advantage of
being insensitive to poisoning by impurities, such as arsenic, anti-
mony, selenium, and chlorine, that may be present in the sulfur.
Since, technically, the greatest cost in processes that depend upon
a catalyst is keeping the catalyst entirely free from minute traces of
poisons, a catalyst less affected by such impurities is preferred. The
introduction and use of the vanadium catalyst was soon followed by
a platinum catalyst that is immune to arsenic poisoning and may
be operated at a lower temperature.
The change of sulfur dioxide to sulfur trioxide, which occurs on
the surface of the catalyst, is an exothermic one, that is, heat is given
off. Careful temperature control is necessary to insure highest yields.
Below 450° C., some SO2 escapes oxidation to SO3. Above 450 °C.,
some SO3 is decomposed to SO2. At a temperature of 450°C. the
reversible reaction
2SO2 + O2 *=± 2SO3
goes to the right almost completely. Only about three percent of the
total amount of SO3 theoretically possible is changed back to SO2.
3) The absorption towers. From the converter, the sultur trioxide
passes through a cooler, and then into a series of three or four ab-
sorption towers. Through the first two of these towers, filled with
quartz pebbles or acid-resisting packing rings, concentrated sulfuric
acid is circulated. This concentrated sulfuric acid absorbs the sulfur
trioxide, forming sulfuric acid of more than 98-percent concentra-
tion. SO3 vapor is absorbed more easily by concentrated H2SO4 than
by water. The third tower is a coke-filled filter tower, which absorbs
any acid vapor that might otherwise escape into the air.
A contact plant is practically automatic. Control of the process is
maintained from one central room containing all the recording and
controlling instruments.
How the contact process compares with the lead-chamber process
of preparing sulfuric acid. The older commercial process for making
sulfuric acid is the lead-chamber process. It was introduced in Bir-
mingham, England, in 1746 by Dr. Roebuck, a physician. Its use
SULFURIC ACID
311
marked the beginning of chemical manufacture on a large scale,
for the Industrial Revolution was just then getting under way.
In recent years, the lead-chamber process has given way gradually
to the contact process. Because of the large volumes of gases that
must react (without a metallic catalyst) , large reaction chambers
are needed. On account of the action of dilute su If uric acid on iron,
it is necessary to line these large rooms with lead sheets, two inches
thick — a second item of expense. Some of the chemical reactions
that occur in the chamber process are still not thoroughly under-
stood.
Sulfur dioxide, steam, and the oxygen of the air are converted into
dilute sulfuric acid by means of the oxides of nitrogen, which act as
catalytic agents. The change of NO into NO2, which takes place in
air, and the subsequent reversal of this reaction are indirect means
of getting an extra atom of oxygen to combine with sulfur dioxide.
At the end of the process a corrosive, sour drizzle of dilute sulfuric
acid falls on the floor of the lead chambers.
Another disadvantage of the lead-chamber process is that only
dilute sulfuric acid, often impure and of not more than 78-percent
concentration, can be made. If concentrated acid is required, this
dilute acid must be concentrated by evaporation with heat.
The chamber process can compete with the contact process only
if acid of not greater than 78-percent concentration is required, as
in the manufacture of phosphate fertilizers and in the pickling of
steel. Sulfuric acid of less than 78-percent concentration must be
shipped in glass or in lead-lined steel containers, while acid of
In these towers the final steps of the contact process take place.
Monsanto Chemical Company and Garfield Chemical and Manufacturing Corporation
312
NEW WORLD OF CHEMISTRY
Fig. 70. Diluting concentrated
sulfuric acid. Why is such care
not necessary when diluting hy-
drochloric acid?
SULFURIC ACID
313
H20
H2S04
greater than 78-percent concentration may be shipped in steel drums
or tank cars. Why?
Physical properties of sulfuric acid. Sulfuric acid, or oil of vitriol,
is a water solution of hydrogen sulfate (a liquid) . Concentrated sul-
furic acid contains 98 percent acid, has a specific gravity of 1.84, and
boils at 338°C. It mixes with water in all proportions, liberating
much heat. For this reason, great care should be used in diluting
concentrated H,SO4. The heavier sulfuric acid should be added to
cold water slowly, and the mixture constantly stirred as shown in
Fig. 70. Always add the acid to the water. Oleum, or fuming sulfuric
acid, contains dissolved sulfur trioxide, and usually its formula is
written H,SO4 • SO3.
Chemical properties of sulfuric acid. A water solution of sulfuric
acid is a very strong acid, because it dissociates to a high degree. The
concentrated acid, because it does not dissociate to any appreciable
extent, is weaker than the dilute acid.
Because of its strong attraction for water, sulfuric acid is an excel-
lent drying, or dehydrating, agent. Its attraction for water is so
intense that sulfuric acid will remove the hydrogen and oxygen pres-
ent in carbohydrates in which these elements occur in the same pro-
portion as in water. Thus, when concentrated sulfuric acid is poured
over sugar, C12H,2On, it removes the H,2OU as 1 1 molecules of steam,
and a black mass of pure carbon remains.
-» 11H2O + 12C
Likewise, wood, which is composed of cellulose, C(1H10O5, chars
when clipped in concentrated sulfuric acid. Cotton, wool, and other
substances react similarly. The dehydrating action of sulfuric acid
accounts also for the severe burns it may produce on the skin.
At 450°C. and atmospheric pressure, sulfuric acid decomposes
completely into water and sulfur trioxide. This chemical reaction
is reversible.
H2SO4 ^ SO3 + H2O
It has been pointed out that nitric: acid is an oxidizing agent be-
cause, when added to a metal, it forms water instead of hydrogen,
and liberates a gaseous oxide. Warm concentrated sulfuric acid,
which behaves in a manner similar to nitric acid, is also an excellent
oxidizing agent.
Cu + 2H2S04 -> CuS04 + 2H2O + SO2 1
This reaction between concentrated sulfuric acid and copper illus-
trates another common laboratory method for preparing sulfur di-
oxide. In contact with certain metals, including zinc, iron, and
magnesium, dilute sulfuric acid liberates hydrogen as shown in the
following equation,
Mg + H2SO4 -» MgSO4 +H2 1
Iron and steel (pickling)
Rayon and film
Other metallurgical
Other
314 NEW WORLD OF CHEMISTRY
Sulfuric acid aids agriculture. In a recent year, the United States
produced about 50 percent of the sulfuric acid consumed by the
entire world. The domestic consumption amounted to slightly more
than 12 million tons. While it is not likely that most farmers believe
themselves in any large measure dependent on sulfuric acid, it is a
fact that the fertilizers used very widely in agriculture account for
almost one-fourth of the sulfuric acid consumed each year.
Phosphorus occurs in nature in a fairly plentiful mineral called
tricalcium phosphate, or rock phosphate. Unfortunately, rock phos-
phate is insoluble in water and therefore cannot be utilized by
plants. However, when treated with H2SO4, rock phosphate changes
into a soluble compound monocaldum phosphate (the active ingre-
dient of superphosphate — see page 487) . The phosphorus is then
available as plant food.
In cases where the soil is acid, however, pulverized raw rock phos-
phate is applied directly. The acids of the soil attack the phosphate
slowly and make part of it available as plant food.
Sulfuric acid in production of petroleum products. The petroleum
industries consume about one-eighth of all the sulfuric acid produced
in the United States. Crude petroleum contains a large number of
carbon compounds that are dark in color or that become dark from
exposure to air. These compounds are removed from crude oil by
treatment with sulfuric acid. In this way most of the waxy and
gummy materials that clog burners and carburetors are removed.
Other uses of sulfuric acid. Another important use of sulfuric acid
is in the cleaning, or pickling, of sheet steel before covering the steel
with a layer of zinc in galvanizing, or tin in tin-plating. The iron
sulfate that is formed as a byproduct is crystallized and is used in
making inks. The common name for crystallized iron sulfate is
copperas, once known as green vitriol.
Fe + H2SO4 + 7H2O -» H2 1 + FcSO4 • 7H2O (copperas)
The manufacture of explosives is dependent upon large quanti-
ties of concentrated sulfuric acid. It is used as a dehydrating agent
in the process of nitrating the many substances from which explo-
sives are made. Sulfuric acid is used in the production of textiles and
rayon and cellulose film. Many acids, including hydrochloric and
nitric, are made with the aid of sulfuric acid. Others of the hundreds
of processes in which sulfuric acid is used include the making of sul-
fates, paper, rayon, leather, celluloid, dyes, dn$$p, and paints. (Many
of these processes will be described later.) Sutewj&acid is used also
as the electrolyte in lead storage batteries. ^%r
SULFURIC ACID 315
Normal, acid, and basic salts. Zinc sulfate, ZnSO4, is formed by the
reaction of zinc and sulfuric acid. Here both of the hydrogen ions
of the acid have been replaced by the metal, and the resulting salt
is a normal salt, that is, a salt in which all the hydrogen of the acid
has been replaced by a metal. In addition to forming normal salts,
acids such as H2SO4, H2COS, and H2SO3, which contain two hydro-
gen atoms, can also form acid salts, and are called dibasic acids. Boric
acid, H3BO3, which contains three replaceable hydrogen atoms, is a
tribasic acid.
In the preparation of hydrochloric acid by the action of sulfuric
acid on sodium chloride, sodium bisulfate, NaHSO4 (also called so-
dium acid sulfate) , is obtained as one of the byproducts.
NaCl + H2SO4 -> NaHSO4 + HC1
Here only one of the hydrogen ions has been replaced, and the other
remains in the salt, which is called an acid salt. Thus it can be seen
that either one or both of the hydrogen ions of sulfuric acid may be
replaced by a metal.
Acids such as HC1 and HNOa contain only one hydrogen atom and
are called monobasic acids. A salt containing one or more hydroxide
groups is called a basic salt. Such a salt is basic lead carbonate
Pb8 (OH) 2 (CO,) a (see page 449) .
Glauber salt and epsom salt. One of the most famous chemists of
the seventeenth century was a Bavarian, Johann Glauber. When he
was 21 he was attacked by a fever and was advised to drink the water
of a certain well. After recovering from his illness he analyzed wa-
ter from this well, and extracted from it crystals of a salt which he
called sal mirabile, the wonderful salt. He recommended it as a
"splendid medicine for internal and external use." This salt,
Na2SO4. 10H2O, known as glauber salt has been used as a laxative
for more than 300 years. Anhydrous Na2SO4, also known as salt cake,
is used in the manufacture of paper and glass.
Soon after the introduction of glauber salt, an English physician,
Nehemiah Grew, who discovered that plants have sex, extracted
a salt from some springs in the village of Epsom, near London,
England. He wrote a book on the medicinal value of this salt,
MgSO4 • 7H2O, and thus epsom salt, a white, soluble, crystalline
compound became a rival laxative.
Many proprietary medicines, mineral-water crystals, and mineral
waters sold today at ridiculously high prices are laxatives containing
one or both of these salts. Such salt mixtures and solutions are sold
to be used, often dangerously, as "the modern way to slenderize."
316 NEW WORLD OF CHEMISTRY
Reducing preparations of this kind should not be taken without the
advice of a competent physician.
How to test for a sulfate ion. Most sulfates are soluble in water.
However, the sulfates of barium, lead, and strontium are insoluble.
When barium chloride is added to a sulfate, a white precipitate,
barium sulfate, is formed. This white precipitate is distinguished
from other barium salts by the fact that barium sulfate is insoluble
in hydrochloric acid.
Na2SO4 + BaCl2 -» BaSO4 J + 2NaCl
This reaction goes to completion. Why?
YOU WILL ENJOY READING
Aaron, Harold. Good Health and Bad Medicine. Consumers
Union of United States, Inc., New York, 1940. Carefully pre-
pared and well-written materials on self-medication.
Goldblatt, L. A., Ed. Collateral Readings in Inorganic Chem-
istry. D. Appleton-Century Co., New York, 2nd series, 1942.
Article 15 contains an excellent description of a "Contact
Sulfuric Acid Plant."
Leicester, H. M., and Klickstein, H. S. Source Book in Chem-
istry. McGraw-Hill Book Company, New York, 1952. Pages
11-16 contain Georg Agricola's description of the manufacture
of vitriol, from De re Metallica.
USING WHAT YOU HAVE LEARNED
Group A
1. Copy and complete the following statements. Do not
write in this book. SO2 is prepared for industrial use by burn-
ing ... or roasting .... It occurs in nature in small amounts
near .... The most important use of SO2 is in the manufacture
of .... Its most characteristic physical property is ....
2. (a) What two methods are used in preparing SO2 in
the laboratory? (b) Which method involves reduction? (c) Ex-
plain.
3. (a) Make a labeled diagram showing how SO2 is pre-
pared from Na2SO3. (b) Write the equation for the reaction
that occurs.
4. Compare the water solubility of SOa,*faH3, HC1, and O2.
SULFURIC ACID 317
5. Write three equations illustrating three chemical prop-
erties of SO2.
6. Describe three uses of SO
2.
7. For what purpose is SO2 used in the household?
8. (a) Is SO2 formed first by the action of H2SO4 on
Na2SO3? (b) Explain.
9. (a) Write the reversible reaction between SO2 and
H2O. (b) How can the reaction be made to go to the right?
(c) to the left?
10. (a) Why is H2SO3 a weak acid? (b) What is a sulfite?
1 1 . Write an equation showing the formation of Na2SO3.
12. (a) By what chemical process does H2SO3 bleach?
(b) Into what does H2SOS change as it bleaches? (c) Compare
the way in which H2SO3 bleaches with the way in which C12
bleaches.
13. Write H2SO3, H2O2, and C12 in a column. Opposite each
write the names of the substances that it is best suited to
bleach.
14- (a) Why is bleaching with H2SO3 less permanent than
bleaching with other chemicals? (b) Illustrate.
15. Distinguish between a sulfite, a sulfate, and a sulfide.
16. How much S would be needed to prepare 500 tons of
98 percent H2SO4?
17. Cu is added to concentrated H2SO4 and 448 ml. of SO2
are liberated. How much H2SO4 is decomposed?
18. Calculate the percentage of S in each of its two oxides.
19. What are three properties of SO3?
20. Write the equation for the burning (roasting) of FeS2.
21. What are the stages in the contact process of making
H2S04?
22. Write equations that show the three chief chemical
changes that take place in the preparation of H2SO4 by the
contact process.
23. (a) Why must the catalyst used in the contact process
be chosen with special care? (b) What advantage has V2O5
over the old plantinum catalyst used in the manufacture of
H2S04?
318 NEW WORLD OF CHEMISTRY
24. Describe briefly the chemical change that is the basis
of the chamber process for making H2SO4.
25. (a) What are the disadvantages of manufacturing
H2SO4 by the chamber process? (b) This process competes
favorably with the contact process in producing H2SO4 for
what industry? (c) Why?
". i " " " ;
26. What are the physical properties of oil of vitriol?
27. In mixing concentrated H2SO4 and H2O, why should
the acid always be added to the H2O rather than the H2O to
the acid?
28. Devise an experiment to show the dehydrating action of
H2S04.
29. By its action on Cu, show that concentrated H2SO4 is
an oxidizing agent.
30. (a) Make a list of the uses of H2SO4. (b) From the
viewpoint of the amount consumed, what is its chief use?
(c) Explain fully, using an equation to show the reaction in-
volved.
31. Why does the refining of petroleum require immense
quantities of H2SO4?
32. Name three experiments in which you used H2SO4 in
making laboratory preparations.
33. What is sulfite pulp?
34. What is the difference between a normal, a basic, and an
acid salt?
35. (a) Name three sulfates and (b) state an important use
of each.
Group B
36. Why is SO3 not added directly to water in making
H2SO4 in the last stage of the contact process?
37. Make a diagram of the apparatus you would use in pre-
paring SO2 in the laboratory by the action of H2SO4 on Cu.
38. Explain: (1) Black rings often form on wooden shelves
that hold bottles of H2SO4. (2) A full battle of concentrated
H2SO4 overflows when exposed to air. (3) An open bottle of
concentrated H2SO4 is sometimes placed in a desiccator.
(4) Frightful burns result from getting concentrated H2SO4 on
the hands. *
SULFURIC ACID 319
39. Compare the three strong acids with respect to their
action on Cu and Zn. Consider both dilute and concentrated
acids.
40. A load of scrap iron weighing 2715 pounds and con-
taining 95.5% iron is added to a large vat of dilute sulfuric
acid. What is the maximum amount of ferrous sulfate that can
be obtained?
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. If you happen to have epsom salt in your medicine chest,
bring it to school, read the label to your class and make a test
for a sulfate with some barium chloride solution your teacher
will supply.
2. Write a two- or three-page report to justify the statement
that H2SO4 is "king of chemicals," and the "basis of the
machine age." Prepare a debate or a class discussion based on
these statements.
22
ALLOTROPIC CARBON:
KEY SOURCE OF ENERGY
. . . It (the light from my lamp)
was hurled out of the sun 200,000,-
ooo years ago, and was captured by
the leaves of the Carboniferous tree-
fern forests, fell with the falling
plant, was buried, fossilized, dug up
and resurrected. It is the same light.
And, in my little fig tree as in the
ancient ferns, it is the same unchang-
ing green stuff from age to age.
Donald Culross Peattie in Flowering
Earth, 1939.
The fascinating chemistry of carbon. Carbon is the eleventh most
abundant element of the earth by weight. In the form ot coal and
graphite, it helps to drive as well as lubricate the wheels of our
machine age. As petroleum, it propels our locomotives, ships, auto-
mobiles, and airplanes. In the form of diamond and "Carborundum"
(synthesized directly from free carbon) , carbon is used in making
tools that drill our hardest rocks and grind and polish our machines,
tools, and household utensils.
In combined carbon we find one of the most fascinating stories
in the whole romance of chemistry. Each year carbon's 650,000
known compounds are being increased in number by thousands of
compounds synthesized in our industrial and university laborato-
ries. Synthetic chemistry reaches its greatest development in organic
chemistry, the chemistry of the compounds ot carbon.
Woehler (vu-ler) , in 1828, synthesized urea, the first organic com-
pound made outside the living body. By this synthesis the old idea
that organic compounds could be formed only in living matter was
destroyed, and a new era in chemistry was ushered in, witnessing
"the great tragedy of science, the slaying of a beautiful hypothesis
by an ugly fact."
320
ALLOTROPIC CARBON
321
One hundred and twenty-five years after Woehler's achievement,
sugar (C^H^O^) was synthesized from its elements. What a century
and a quarter of research between Woehler's urea and synthetic
cane sugar! Berthollet went to the red ant and learned the secret of
preparing formic acid, the liquid responsible for the sting of this in-
sect. William Perkin, washing bottles in a London laboratory, mixed
at random the contents of two bottles and discovered a method of
synthesizing mauve — the first of a long series of organic dyes that
rival the colors of nature. And this is only the beginning. The mind
fairly reels at the accomplishments and possibilities of this branch of
chemistry. More of this entrancing story is told in later chapters.
What is the valence of carbon? Carbon has an atomic weight of
12. The nucleus of the carbon atom contains six protons; hence, the
outer ring contains only four electrons. Therefore, the carbon atom
must either lose four electrons or gain four electrons to complete its
outer ring. The valence of carbon then is either plus four or minus
four, depending upon the substance with which it reacts. That is,
it may act as either a metal or a nonmetal. Carbon almost always
forms firm covalent bonds involving four pairs of shared electrons.
The properties of carbon. Carbon occurs in three allotropic forms:
two of them, diamond and graphite, are crystalline; the third is
amorphous, or noncrystalline, in form. All forms of carbon combine
with oxygen, forming either carbon monoxide, CO, or carbon di-
oxide, CO2. The different forms of carbon vary widely in the ease
with which they may be oxidized. Coal burns easily particularly
when powdered. On the other hand, diamond combines with oxygen
only at extremely high temperatures.
C + O2 — > CO2 (carbon dioxide)
2C + O2 — > 2CO (carbon monoxide)
Its attraction for oxygen makes carbon a good reducing agent. The
reducing action of carbon may be shown by heating a mixture of
blowpipe
Carbon
Fig. 71. Blowpipe analysis
of an ore (left) using the re-
ducing power of charcoal,
(right) The carbon atom.
322 NEW WORLD OF CHEMISTRY
carbon and lead oxide on a charcoal block, using a blowpipe. A lit-
tle ball of pure lustrous lead is formed and CO2 is given off.
2C 4- Pb3O4 -> 3Pb + 2CO2 T
Millions of tons of coke (carbon) are used industrially every year
in reducing iron oxide to commercial forms of iron.
Fc2O3 + 3C -> 2Fc + SCO |
This reaction shows carbon acting as a metal, a lender of electrons.
Carbon may also behave as a nonmetal, a borrower of electrons. At
the very high temperatures reached in an electric furnace, carbon
unites with metals such as calcium and wolfram and with less electro-
negative elements such as silicon, forming carbides.
CaO +3C-> CaC2 + CO t
calcium oxide calcium carbide
SiO2 +3C-» SiC +2COT
silicon dioxide silicon carbide
Carbon unites directly with hydrogen also at the high temperatures
of an electric arc, forming acetylene gas, C..H.,.
2C + H2 -> C2H2 T
All forms of carbon are insoluble in water, acids, and bases.
Amorphous carbon, however, is soluble in some molten metals in-
cluding iron.
Physically, the allotropic forms of carbon resemble one another
in some respects. They are all odorless, tasteless, and insoluble in
water. On heating, they do not melt but vaporize.
The hardest substance known. In the eighteenth century, diamond
was shown by Lavoisier to be almost pure carbon. This French chem-
ist burned a diamond in the presence of a distinguished audience by
concentrating the sun's heat on it with a lens. He showed that the
gas which was formed was pure carbon dioxide. This startling dis-
covery strengthened Lavoisier's theory of burning. Why?
Diamond, chief among precious stones, is found mainly in South
Africa, Brazil, and Borneo. It occurs in various sizes and states of
purity, buried deep in the ground frequently in the pipes of extinct
volcanoes, or in the loose sands along certain streams. The brilliance
of diamond, caused by its power of refracting light, is brought out
by removing any surface impurities or imperfectiqns and by cutting
facets on its surface. Small amounts of metallic oxides in diamonds
ALLOTROPIC CARBON
323
Shell Oil Company Baunigold Bros., Inc.
This asphalt core, neatly cut from a Before use in jewelry, a diamond is
roadbed by a diamond drill, will be polished against a rotating iron disc
taken to a laboratory for study. coated with olive oil and diamond dust.
give them various colors, the rarest of which are red, green, and blue.
The density of diamond is 3.5.
Diamond is weighed in carats, a unit of weight based on the weight
of the bean of the carob tree. One carat is equal to one-fifth of a
gram. The largest diamond ever mined was the flawless Cullinan
found in South Africa in 1905. It weighed 3026 carats, or 21.35
ounces, and was about the size of a man's fist.
From the first little diamond found in a pebble picked up by a
child in 1867, the diamond industry of South Africa has grown to
world leadership. In a recent year, more than three million carats of
diamonds were mined. About half of this amount is used in jewelry.
Though brittle, diamond is the hardest substance known, and for
this reason is used extensively in making tools for drilling rock,
grinding, sawing, engraving, and polishing/ Impure black diamonds,
called carbonado and bort, are used in drills.) Diamonds are also used
as bearings in watches and delicate instruments, and in wire-drawing
dies made by drilling tapered holes through diamond crystals.
Can diamonds be made synthetically? Scientists believe that
diamonds were formed under great heat and pressure by the crystal-
lization of carbon dissolved in molten iron or other substances. To
test this theory of the formation of diamonds, Henri Moissan (mwa'-
sa'N') , in 1893, subjected some pure carbon to extremely high tem-
peratures and pressures. He dissolved carbon in molten iron, and
quickly plunged the molten mass into cold water. The sudden
contraction of the iron exerted tremendous pressure on the dissolved
324 NEW WORLD OF CHEMISTRY
carbon, causing it to crystallize. The iron was dissolved with acids,
and, aided by a microscope, Moissan reported seeing synthetic dia-
monds formed of tiny cubical crystals. Other attempts have been
made to synthesize diamonds, but thus far these attempts have not
proved a threat to the diamond industry.
Black, soft, and greasy. A second allotropic form of carbon,
graphite, is lustrous, black, soft, and oily to the touch. The great
difference in hardness between the allotropic forms of carbon is
caused by the different arrangement of the atoms. The atoms of the
diamond molecule are very closely interlaced and bound firmly to-
gether. Graphite, on the other hand, is made up of piles of flat lay-
ers of atoms firmly bound to one another, but far apart and only
feebly hanging on to the next layer. For this reason graphite is soft
and slippery. The flat planes slide over one another like cards in a
pack. A few electrons between the planes are loosely held, making
graphite an excellent conductor, whereas diamond is a nonconductor.
Natural graphite, which is found chiefly in Ceylon, Siberia,
and Canada, contains many impurities. It is believed to have been
formed, like diamond, under high temperature and pressure.
Electric furnaces and uses of synthetic graphite. In 1896 Edward
Acheson, ot Pennsylvania, patented a process for making synthetic
graphite,, today a widely used article of commerce. It is superior to
natural graphite, because it is free from grit and contains almost no
impurities. Acheson was at different times a blast-furnace helper, oil
gager, raih*oad ticket agent, store clerk, miner, bookkeeper, and en-
gineer. He also worked for Edison and was sent abroad to the Inter-
national Exposition in Paris to exhibit the Edison electric lamp
which he had helped to improve with a graphite filament. His
method of making synthetic graphite consists of subjecting anthra-
cite coal or molded carbon to the high temperature of an electric
furnace, keeping air out by covering the coal with sand.
An electric furnace, as you know, transforms electric energy into
heat energy. The resistance to an electric current offered by loose
carbon, in the center ot the resistance type of electric furnace used
in making graphite, produces high temperatures. Other electric fur-
naces utilize the heat produced by the passage of electricity through
carbon atoms
Fig. 72. Diagram of a graphite
particle showing the layers of
alom,.
crystals
Alloy steels are produced
in this electric furnace. The
electrodes of the furnace
are made from graphite.
Electro Metallurgical L'o mint try
high-resistance wire, such as an alloy of nickel and chromium. An-
other type of electric furnace is the arc type. Here the passage of
electricity across an air gap produces a high temperature as a result
of the resistance of the air. Arc-type electric furnaces are used in the
synthesis of nitric oxide, NO, during the preparation of nitric acid
by nitrogen-fixation.
Like other forms of carbon, graphite does not melt and has an ex-
tremely high point of vaporization. For this reason it is used in mak-
ing graphite crucibles in which steel and other alloys of high melting
point are prepared. Because of its smoothness, oilincss, and high tem-
perature of fusion, graphite is mixed with oils used in lubricating
heavy, swiftly moving parts of machinery, in which even heavy oil
would evaporate or burn off.
Graphite electrodes are used in electric furnaces, arc lamps, and
in general, for conducting electricity at very high temperatures.
Powdered graphite is used in dry cells, and in certain oil-retaining
bearings produced by powder metallurgy (see page 467) .
Graphite is used also in lead pencils. The "lead" in a pencil is
graphite mixed with clay to vary the hardness of the lead. Graphite
is also a constituent of stove polishes. Most synthetic graphite is pro-
duced at Niagara Falls, where hydroelectric power is available at
low cost.
sand
C0dlv :| y Fig 73 Cross section of an
electric furnace used in the
commercial production of
synthetic graphite.
graphite electrodes
325
326 NEW WORLD OF CHEMISTRY
"Formless," yet of many forms. The noncrystalline allotropic form
of carbon is amorphous carbon. The purest form of amorphous car-
bon can be prepared easily by heating pure table sugar. The hydro-
gen and oxygen are driven off as water; pure amorphous carbon is
left.
Ci2H22Oii -> 11H2O + 12C
Among the most common forms of amorphous carbon are (a) lamp-
black and carbon black, (b) gas carbon, (c) coke, (d) charcoal, and
(e) animal charcoal, or boneblack.
Valuable soot. Lampblack was known to the Chinese and Egyp-
tians who used it in making ink. It is amorphous carbon dust formed
by the incomplete burning of compounds rich in carbon. It is man-
ufactured commercially by burning heavy liquid hydrocarbons (com-
pounds of hydrogen and carbon) in an insufficient supply of oxy-
gen.vThe escaping particles of carbon dust, or soot, from the small
flames are collected on a revolving metal plate that is kept at a low
temperature by flowing water. The smokiness of a kerosene lamp
burning with an insufficient air supply or with the oil supplied too
rapidly illustrates the way in which lampblack is formed.
Carbon black is made in the same way as lampblack except that
natural gas is used instead of liquid hydrocarbons. It* is finer than
lampblack, which it is gradually replacing. Carbon black is such an
important item of commerce that in a recent year its production ac-
counted for more than ten percent of all the natural gas produced in
the United States. Most of the carbon black manufactured in this
country is added to rubber tires to improve their resistance to tear
and abrasion. It is also used as a basic raw material in making print-
er's ink, typewriter ribbons, carbon paper, phonograph records, and
black paints.
Almost graphite. Gas carbon is amorphous carbon that collects on
the walls of retorts during the manufacture of illuminating gas from
coal. In structure, gas carbon is semicrystalline and resembles graph-
ite. It is used in the rods of arc lamps and other electrodes.
Coke from the destructive distillation of coaL Heating a complex
carbon compound in an oven from which air is* excluded, and con-
densing the vapors formed is called destructive distillation. When
soft coal is so treated the most important solid product is coke, a
steel-gray, hard, brittle substance used chiefly as a reducing agent
in the extraction of iron, and as a fuel. The coal is heated to 2000 °F.
for 16 hours, and then the white-hot coke is pushed out from the
oven and quickly quenched in cold water.
ALLOTROPIC CARBON 327
In the byproduct coke oven many valuable chemicals are recov-
ered. As the gases pass through oil-absorption tanks, a light oil sepa-
rates, from which benzene, toluene, xylene, and naphthalene are dis-
tilled. From the coal tar which is formed anthracene, phenol, and
pitch are also separated by fractional distillation. Ammonia gas and
coal gas are also obtained (see pages 362-363 for a detailed account
of this process) . Less than ten percent of our coke is still made in
beehive ovens from which none of these byproducts are recovered.
Charcoal-broiled. If wood is subjected to destructive distillation,
charcoal is one of the products. Charcoal retains the fibrous struc-
ture of wood.
Powdered charcoal is frequently put together under tremendous
pressure, forming compact, pillow-shaped blocks called charcoal
briquets. These burn with no smoke and leave very little ash. They
are used in camping or picnicking, on pleasure boats, and in found-
ries.
Charcoal-broiled meats, especially steaks and chops, are particu-
larly delicious. Meats broiled over charcoal do not pick up odors or
flavors from the burning fuel.
Activated charcoal. This is a specially-prepared charcoal used for
removing or adsorbing gases. Adsorption is caused by the collection
of thin layers of molecules of gases and other impurities on the sur-
faces of the porous charcoal. It is not a chemical union but a physical
attraction. Ill-smelling colored liquids, passed through such charcoal,
are cleared of impurities responsible for the odor and color. Adsorp-
tion is used also in recovering many industrial solvents and waste sub-
stances, in air purification, and to concentrate a desired substance
such as the drug streptomycin from its mold culture.
Activated charcoal is used in gas masks to adsorb poison gases met
in industry. "Nuchar" is the trademark for an activated carbon of
vegetable origin containing about 90 percent pure carbon.
A black decolorizer. The chief solid left after the destructive
distillation of clean cattle bones is a black powder called boneblack.
This boneblack, or animal charcoal, is a mixture containing only
about ten percent amorphous carbon and 80 percent calcium phos-
phate, the chief compound found in bones. It, too, has great adsorp-
tive powers and is used mainly in the refining of crude sugar, from
which it removes impurities that cause objectionable colors and
odors. It is used also as a black pigment in paints.
How coal was formed. Coal is found in all continents, including
Antarctica, but the largest coal-producing areas are in North Amer-
ica and Asia. Scientists now believe that many millions of years ago,
328 NEW WORLD OF CHEMISTRY
during the Carboniferous period, great portions of the earth were
covered with a dense vegetation more luxuriant than that found
today even in tropical jungles. As the level of the earth sank during
one of the many upheavals that occurred, vast portions of these
jungles were flooded by the ocean and became swamps. Later these
swamps were completely submerged and then gradually covered with
mud, sand, or clay deposited as sediment by streams and rivers.
Partial decomposition of the wood and other vegetable matter,
aided by bacteria and fungi, changed the woody material first into
peat. Minnesota, Florida, and several other states contain large peat
deposits. Near Chester, Wisconsin, is a peat deposit that covers
32,000 acres. Its average depth is six feet and it contains about
40,000 tons of air-dried peat. Peat bogs also occur in Ireland and else-
where. When dried and pressed, peat may be used as a fuel. It burns
with a great deal of smoke and produces little heat.
Scientists believe that the next step in the formation of coal was
the changing of the peat into brown or black lignite^ which still re-
tains the structure of the plants from which it was formed. Some-
times branches or twigs are found in lignite in the form in which
they grew. Lignite is found in some of the states west of the Missis-
sippi, and, where no other form of coal is available, is used as a
household fuel. In North Dakota, for example, lignite is being
mined at the present time and utilized in making gas and briquet
fuels, and for the production of electric power.
Further decomposition and pressure, away from air, drove out
more oxygen and hydrogen from the lignite, leaving hydrocarbons
and some free carbon and forming bituminous, or soft, coal. Further
heat and greater pressures changed the soft coal to jet black, lus-
trous hard coal, or anthracite.
Strong evidence supports this theory of coal formation. The fos-
sils embedded in coal deposits include forms of both animal and
vegetable life. Some of the animals appear to have been marine.
Among the plants, mosses and ferns occur most commonly.
The chief kinds of coal. The chief kinds of coal used in this
country are soft, or bituminous, coal, and hard coal, or anthracite.
Most of the anthracite mined in the United States comes from Penn-
sylvania, yet this state, together with West Virginia, produces more
than half the bituminous coal mined in this country. For every ton
of anthracite mined, 11 tons of bituminous coal are dug. Anthra-
cite contains practically no free carbon and produces less volatile
material than bituminous coal (see Table 11) . It is cleaner to han-
dle and burns with a short, pale blue flame.
ALLOTROPIC CARBON 329
The use of anthracite is confined largely to the northeastern states,
because these states have the advantage of being near the deposits
and thus have short and low-cost transportation from the mines. In
other parts of the country anthracite is relatively unfamiliar.
The problem of smog. Smog, a mixture of smoke and fog, is
formed largely by the incomplete combustion of coal, fuel oil, gas
and rubbish. The products of this incomplete combustion are oxi-
dized by the o/one in the air to create the smog. Smog may be aggra-
vated by natural particles in the air, such as dust or pollen, or even
by local weather conditions.
People who live in cities where smog is a problem spend far more
for soaps and other cleansers than might be necessary if better smoke
control were practiced. It has been shown in several cities that smog
may be eliminated. However, success in reducing smoke depends
upon the cooperation of every consumer of fuel from the largest
industry to the individual homeowner.
Smoke control is particularly necessary in cities in which bitumi-
nous coal is a major fuel. Bituminous coal, when burned in an in-
adequate supply of oxygen, produces a very smoky flame. If the coal
is properly fired (sec page 35!)) , smoke may be greatly reduced.
British thermal units. A chief factor in determining the price of
coal is its fuel value. This value is measured either in calories or in
British thermal units, Btu. One Btu is the amount of heat required
to raise the temperature of one pound of water one degree Fahren-
heit. One Btu is equivalent to 252 calories. Burning one pound of
average anthracite liberates about 12,700 Btu; burning the same
weight of bituminous coal liberates about 13,100 Btu. The per-
centage of ash and free sulfur is another factor important in deter-
mining price. In general, anthracite contains a smaller percentage
of these materials than does bituminous coal.
A scene in Pittsburgh, Pa., before (left) and after a smog-control program was
instituted.
Allegheny Conference on Community Development
TABLE 11. fl
VARIATIONS ||
IN CHEMICAL f|
CONTENT OF COAL 1
Variations in fixed ||
carbon, volatile matter, ^
and moisture on an ash ^
free basis of the several ||
ranks of coal produced ^
in the United States ||
Moisture
Lignite
Subbituminous C
Subbituminouj B
Subbituminous A
High -Volatile C Bituminous
High -Volatile B Bituminous
High -Volatile A Bituminous
Medium -Volatile Bituminous
Low -Volatile Bituminous
Semianthracite
Anthracite
Meta- anthracite
20%
60*
80*
Coal, a major source of power. Coal has been used as fuel for
centuries. Marco Polo, 700 years ago, reported the Chinese burning
"black stone dug out of mountains where it runs in veins." But it was
not widely used until the sixteenth century. By 1661 "a hellish and
dismall cloud of Sea-Coale hung perpetually over London," where
coal was used to make iron, glass, and other products. During the
Industrial Revolution it became a vital source of energy.
Our civilization uses energy at a staggering rate. In 1900, the total
supply of energy from mineral fuels and water power in the United
States was 7.4 quadrillion Btu; by 1950, the total had reached 36.2
quadrillion Btu — an increase of almost 400 percent! Each year our
growing population and industries need added amounts of energy.
It is expected that before the end of the century the annual produc-
tion of energy will be more than double the present rate.
Today, petroleum, coal, natural gas, and hydroelectric dams are
our major sources of energy. In a recent year, petroleum supplied 40
percent of our nation's total energy requirement; coal supplied 34
percent; natural gas, 22 percent; and water power, 4 percent. In
the years ahead, the use of nuclear energy will, in all probability,
change these figures radically.
Coal has many uses. Production of anthracite coal is now about
40 million tons per year. About three-fourths of this, output goes to
the retail market to be used as fuel for commercial and institutional
buildings, as well as for heating homes. The remainder of the yearly
tonnage is used by a variety of industrial consumers.
Annual production of bituminous coal is now about 450 million
tons per year. The major use of this coal is in generating electricity.
330
ALLOTROPIC CARBON
331
About one-half of our nation's total electric power is produced by
coal-fired steam plants. In all, the electric utilities consume consider-
ably more than 100 million tons of coal each year, or about 25 per-
cent of all bituminous coal marketed. The production of electric
power is expanding rapidly and it is expected that there will be an
increase in the amount of coal used for this purpose.
The steel industry is the second most important customer for
bituminous coal, consuming about 20 percent of all coal marketed.
Most of this enormous quantity of coal is converted into coke by
roasting in coke ovens. Coke is used to some extent as a fuel, but its
chief use by far is as a reducing agent in the production of steel.
Coke robs the iron ore, an oxide of iron, of its oxygen and leaves pig
iron which is later converted to steel. The chemicals recovered in
the coking process (see page 327) are the source of many plastics,
dyes, drugs, medicines, and industrial chemicals. Today, coal chemi-
cals are also produced directly from the "raw" coal without any at-
tempt at burning it or changing it to coke. This use of coal as yet
constitutes a very small market, but it may become increasingly im-
portant in the future. Coal is also destructively distilled in the manu-
facture of coal gas.
All other types of manufacturing consume another 15 to 20 per-
cent of the bituminous coal marketed annually. A similar amount
goes to the retail market for use as a fuel in public buildings, com-
mercial establishments, and homes. In addition to coal in its conven-
tional lump form, the retail market each year uses over a million
tons of coal in the form of fuel briquets. Fuel briquets are pressed
cubes, cylinders, or ovoids of very fine bituminous coal (slack) or lig-
nite held together by a burnable binder.
The railroads, once the coal industry's best customer, now use
considerably less than ten percent of the total annual tonnage as the
Consumers Power Company
The importance of coal in
generating electricity is illus-
trated by this coal storage
pile at a modern power
station. This plant burns
about 675,000 tons of coal
each year.
332 NEW WORLD OF CHEMISTRY
coal-burning "iron horse" has been displaced by the more efficient
diesel-electric locomotive.
Buying coal for your home. The most important consideration is
the amount of heat the coal will produce. Your coal dealer usually
knows the Btu value of the coal. Such information may also be ob-
tained from the United States Bureau of Mines, from your state bu-
reau of mines, or from the company that mined the coal.
If you know the number of Btu in a pound of coal, the product of
this number multiplied by 2000 is the number of Btu per ton. The
quotient of the number of Btu per ton divided by the cost per ton in
dollars is the number of Btu per dollar. As far as heat value alone is
concerned, the coal with the greatest number of Btu per dollar is the
best buy. Expressed as a formula, this method of comparing coals
may be written:
Btu per Ib. X 2000 _ , „
— - : — rr; = Btu per dollar
Cost per ton in dollars r
However, as anyone who has ever carried out "ashes" knows, the ash
content of coal is an important consideration also. One coal with a
slightly higher Btu-per-dollar value than a second coal actually may
not be as desirable as the second coal, if its ash content is appreci-
ably higher. Moisture content and dust content are two other fac-
tors that must be considered. Coal dealers may control the dustiness
of coal either by sprinkling the coal with water or by spraying the
coal with special oils. From the point of view of the consumer, which
of these methods is preferable? Why?
In buying coal, it is advisable to buy the size and grade recom-
mended by the manufacturer of the equipment in which you intend
to burn it. Buying the appropriate size and grade of coal will enable
you to get more Btu out of the coal you buy. Most coal-burning
equipment is designed to operate most efficiently with coal of cer-
tain size and grade.
Because of seasonal demands, coal is usually lower in cost during
summer than in winter. Since coal does not deteriorate appreciably
in dry storage, for example, in a coal bin, it is good policy to buy
next winter's supply in the summertime.
Changing energy relationships. The sources of energy upon which
our civilization depends are, as we know, coal, petroleum, natural
gas, and water power. Nuclear energy holds great promise, but its
use as a major source of energy is still in the future.
Petroleum and natural gas constitute but a tiny part of our na-
tion's total fuel deposits. Taken together, they m^te up less than
ALLOTROPIC CARBON 333
two percent of the known reserves of all fuels. Of the remaining 98
percent, oil shale forms eight percent and coal forms 90 percent.
More than 92 percent of the total coal reserve is bituminous and
lignite; the remainder is anthracite. Although the experts disagree
on the exact number of years involved, it is generally accepted that,
at our present rate of use, our reserves of high-quality petroleum and
natural gas may be consumed in less than half a century, while our
coal reserves are adequate for several hundred years. Oil shale is an
unexploited resource and we do not yet know the extent of its use-
fulness.
Because of coal's abundance, it seems likely that it will see in-
creased use in the future as oil and gas reserves shrink. Because our
society is completely dependent upon a plentiful and continuous sup-
ply of energy, proper use and conservation of our fuel reserves is a
problem which deserves the thoughtful attention of every citizen.
The energy of running water, while limited in amount, is never-
theless almost limitless in duration. For, in harnessing such energy,
nothing is actually used up that will not occur again. The water cycle
takes care of this problem for us. However, water power suffers from
a serious drawback in that it is not constant. That is, during certain
seasons of the year, more water is available than at others. Even
dams and huge reservoirs have not yet succeeded in solving this prob-
lem completely, but they have made tremendous strides in doing so,
and more and more water power is being harnessed as the years go
by. However, even if developed to the fullest extent, our water re-
sources could not provide more than a small fraction of our total
annual energy requirements.
The utilization of hydroelectric power had to wait for science to
develop methods of sending electric energy over long distances, and
also for the building of immense dams and hydroelectric plants.
Recently, the capacity of existing hydroelectric power plants in the
United States totaled 25 million horsepower, while the undeveloped
water power was estimated at 117 million horsepower. Engineers be-
lieve that, long before we can no longer depend upon coal, petro-
leum, natural gas, and the harnessed power of running water, nuclear
energy will do much of the world's work.
Motor fuels from coal. Chemists, faced with an imminent world-
shortage of petroleum, turned to the development of methods of
synthesizing essential liquid fuels. One of the most widely used de-
velopments has been the changing of low-grade coal into gasoline
and other petroleum products. The most successful results were ob-
tained by the German chemists, Friedrich Bergius (ber'gi-obs) , Franz
334
NEW WORLD OF CHEMISTRY
Fischer, and Hans Tropsch. In 1931, Bergius was awarded the Nobel
prize in chemistry.
Bergius found that coal contained half as much hydrogen as petro-
leum. By doubling the amount of hydrogen chemically combined in
coal, he hoped that lie could obtain a product that would resemble
gasoline. He developed a process now called the Bergius process, in
which a thick paste of powdered coal mixed with heavy oil is intro-
duced into steel drums and heated to a temperature of about 400°C.
under about 200 atmospheres pressure. Hydrogen is forced into the
mixture, which also contains a catalyst. When the hydrogenation is
complete, a mixture of gasoline and heavier fuel oils is obtained.
The carbon of the coal combines chemically with hydrogen, form-
ing certain kinds of hydrocarbons, and the successful liquefaction
of coal is achieved.
Assuming the composition of average gasoline to be represented
by the formula C7H1(J, we may represent the hydrogenation of coal
by the equation:
7C+8H2-»C7H16 .-;"::;•:; '
During World War II, the Germans demonstrated that synthetic
gasoline from coal can be used for the same purposes as. petroleum-
derived gasoline. In the United States, following the war> both the
federal government and private industry conducted research on the
production of gasoline from coal and from natural gas. It was found
that synthetic gasoline is more expensive to produce than gasoline
from petroleum, but would be commercially practical in the event
of shortages of the latter.
Apparatus in a coal-industry
research laboratory for testing
stoker coal and thus establish-
ing standards for determining
the best use for the various
grades.
The testing laboratory of a
coal preparation plant. The
crucibles contain powdered
coal which is burned in the
electric furnace. By analysis
of the ash, the composition
of the coal is determined.
National Coal Association
Both the Bergius and Fischer-Tropsch processes were tested. The
Fischer-Tropsch process depends either upon the reaction of steam
and coal, or oxygen and natural gas, to form synthesis gas.
CH-H20->CO + H2 or 2CH4 + O2 -» 2CO + 4H2
synthesis gas synthesis gas
This mixture of gases is then converted with the aid of a catalyst
into gasoline.
iron
7 CO + 15H2 > C7H16 4- 7H2O
• . catalyst ,
synthesis gas synthetic
)line
YOU WILL ENJOY READING
Bragg, William. Concerning the Nature of Things. Dover
Publications, New York, 1954. Lecture IV on crystals covers
the various allotropic forms of carbon. A classic.
Storch, H. H.; Lowry, H. H.; Kiebler, M. W., Jr.; How-
ard, H. C.; Thiessen, Gilbert; and Charmbury, H. B. "Hy-
drogenation of Coal." Industrial and Engineering Chemistry,
April, 1944, pp. 291-298. A round-table discussion of problems
in this process.
USEFUL IDEAS DEVELOPED
1. Destructive distillation is the heating of a complex or-
ganic compound in a chamber from which air is excluded
and condensing the vapors formed.
335
336 NEW WORLD OF CHEMISTRY
2. Fuel value of coal is measured in calories or British
thermal units, Btu. One Btu is the amount of heat required
to raise 1 Ib. of water 1°F. A calory is the amount of heat re-
quired to raise 1 g. of water 1°C. One Btu is equivalent to
252 calories.
3. Organic chemistry is the chemistry of carbon compounds.
At present more than 650,000 compounds of carbon are known.
Synthetic chemistry reaches its greatest development in this
branch of chemistry and has already produced dyes, drugs, and
solvents many of which are better than natural substances.
4. Today, man's chief sources of energy are coal, petroleum
and natural gas, and water power. In the years just ahead,
nuclear energy is likely to assume greater importance.
USING WHAT YOU HAVE LEARNED
Group A
1. C may act either as a metal or as a nonmetal. (a) Why?
(b) Why is C not active at ordinary temperatures?
2. (a) What are the three allotropic forms of C? (b) Name
the allotropic forms of two other elements.
3. (a) What chemical property is common to all allotropic
forms of C? (b) What product depends on this property?
4. What accounts for the extreme hardness of diamond?
5. How did the action of extreme heat on diamond aid
Lavoisier in securing acceptance of his theory of burning?
6. (a) What are two possible products of the oxidation
of C? (b) Write equations to illustrate.
7. (a) What are the three chief uses of diamond? (b) Upon
what properties do these depend?
8. What evidence supports the accepted theory of diamond
formation? ft
9. (a) What are the chief uses of graphite? (b) Upon
what properties do these depend?
10. Describe briefly the commercial preparation of graphite.
11. (a) What is an electric furnace? (b) Name three kinds.
12. Make a diagram of one kind of electric furnace and ex-
plain how it produces high temperatures.
t . .-
13. Make a table of four kinds of amorphous C, giving the
ALLOTROPIC CARBON 337
method of manufacture, composition, and chief uses of each.
14. (a) What is destructive distillation? (b) Describe the
making of coke, and (c) name six byproducts of the process.
15. Activated charcoal is used in gas masks. Why?
16. Below are two lists of items. Match each item in the list
on the left with the correct item from the list on the right.
1) watch bearings a) activated charcoal
2) Moissan b) graphite
3) peat c) printer's ink
4) Bergius d) diamond crystal
5) lampblack e) coke
6) manufacture of iron f) artificial diamonds
7) lead pencils g) bacteria
8) adsorption h) hydrogenation
17. What weight of coke containing 80 percent C is needed
to reduce 5 Ib. of CuO?
T f , T ,
I
18. Assume the formula of gasoline to be C7H16. If 500 ml.
of water vapor are formed during the combustion of a certain
quantity of gasoline, what volume of CO2 will be formed?
19. What volume of CO2 will be formed during the com-
bustion oi; 2000 Ib. of coal containing 80 percent C? (Note:
the ounce-molecular volume of any gas is 22.4 cu. ft.)
20. According to available evidence, how were coal de-
posits formed?
21. Compare anthracite and bituminous coal for use in
household heating.
22. What factors should be considered in buying coal?
23. How may the heat value of one coal be compared with
the heat value of another?
24. Complete and balance the following equations:
a) C + H2 -> c)
b) Fe2O3 + C -> d) SiO2 + C ->
25. What is meant by saying that a certain sample of coal
liberates 14,000 Btu?
26. What is the major use of anthracite coal? What two in-
dustries are the major consumers of bituminous coal?
„ . I .
1
27. Explain the Bergius process for the hydrogenation of
338 NEW WORLD OF CHEMISTRY
coal. Why did he experiment to achieve this result?
28. With the aid of equations explain the synthesis of
gasoline by the Fischer-Tropsch process.
29. (a) What are three physical properties possessed by all
allotropic forms of C? (b) In what physical property do they
differ radically?
30. Explain why the term lead pencil is a misnomer.
31. Give two reasons why graphite is used in making stove
polishes.
Group B
32. 2 g. of coal raised the temperature of 2000 ml. of water
6°C. Determine the fuel value of this coal in calories.
33. Name four reducing agents and give an example of the
reducing power of each.
34. How is an imitation diamond distinguished from the
genuine?
35. Wooden posts are sometimes charred before being placed
in the ground. Explain.
36. Charcoal tablets are sometimes used in the treatment
of certain stomach disorders. Explain.
37. What produces chemical changes in an electric furnace:
the heat or the electric current? Explain.
38. How does a large percentage of ash or S in coal affect
its quality as a fuel?
39. (a) Why is Woehler sometimes called the father of
organic chemistry? (b) What is the difference between organic
and inorganic compounds?
40. How would you identify three black powders in a mix-
ture of graphite, manganese dioxide, and copper oxide?
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Place a small amount ot sugar on a hot stove or hot plate
and report what happens, indicating as many changes as you
observe.
2. Make a small working model of an electric lurnace used
in the production of graphite.
3. Large users of coal purchase coal according to definite
specifications as to the percentage of C, volatile material, ash
and sulfur content, and fuel value. Does your father buy coal
according to definite specifications? Why do most domestic
consumers not bother about specifications in buying coal?
Is this a sensible state of affairs? Explain.
23
CARBON DIOXIDE:
GAS OF LIFE AND DECAY
. . . In /757 Joseph Black discov-
ered that carbonic acid gas could not
be breathed by animals, and had a
poisonous effect on them. Sparrows
introduced into an atmosphere of
the gas died in 10 or n seconds, but
if their nostrils were stopped with
lard, their death took place only at
the end of 3 or 4 minutes. Robert
Routledge in A Popular History of
Science, 1881
A Scottish physician discovers carbon dioxide. Joseph Black, a
physician, was very much interested in a lively discussion between
two professors of medicine. One professor believed that limewater
made from limestone was a more effective medicine than limewater
made from oyster shells. Black, to settle this controversy in true sci-
entific manner, undertook a thorough study of several carbonates.
In 1754, while heating one of these carbonates (magnesium carbo-
nate) , he obtained pure carbon dioxide gas for the first time. This
was 20 years before the isolation of pure oxygen. Black called car-
bon dioxide "fixed air."
MgCO3 -» MgO + CO2 T
The atmosphere is a vast storehouse of this gas. Immense quanti-
ties also occur dissolved in surface and ground water. Large quan-
tities escape from volcanoes and crevices in the earth; it is also
ejected in tremendous volumes from wells that produce petroleum
or natural gas. In fact, enough carbon dioxide could be obtained
from oil wells in the United States alone to supply the world. Carbon
dioxide is formed when carbon or its com pounds are burned. Be-
cause all common fuels contain carbon either free or in compounds,
339
animals
«.. ~ TH. c
oxygen cycle.
green plants in sunlight
carbon dioxide is produced whenever such fuels are burned. Finally,
tremendous volumes of carbon dioxide are locked up chemically in
the great masses of metallic carbonates and bicarbonates found in
the earth's crust.
Why carbon dioxide is necessary for life. Life, as we know it today,
would be impossible without this percentage of carbon dioxide gas
in the air, small though it is (0.04 percent) . For without it, plants
even with the aid of sunlight could not manufacture starch. As a
result, our food supply would diminish and finally vanish altogether.
Since coal was formed in prehistoric days by the destructive distilla-
tion of plants, even this great source of power would never have been
formed if carbon dioxide had not been present in air.
On the other hand, a much larger concentration of carbon dioxide
in the atmosphere would be fatal, for it would dilute the oxygen and
choke the breath of life out of living things. The delicately balanced
ratio of oxygen to carbon dioxide in the air may be better appreciated
after a study of the carbon dioxide-oxygen cycle.
The carbon dioxide-oxygen cycle. Carbon dioxide is constantly
being added to the air by the breathing of animals. It is made in the
body tissues by the oxidation of carbon compounds in these tissues.
Some of the oxygen inhaled in air is changed to carbon dioxide,
which passes out of the body through the lungs. Every breath of air
exhaled contains 100 times as much carbon dioxide as was inhaled.
At the same time, the amount of oxygen decreases from about 2 1 per-
cent in the air inhaled to about 16 percent in the air exhaled.
Every ton of coal and liquid fuel burned sends about four tons of
carbon dioxide into the air, diminishing, at the same time, the oxy-
gen content of the air. During the decay and fermentation of or-
ganic material, immense volumes of carbon dioxide are given off
into the air.
You might suspect that, after hundreds of thousands of years,
these three processes — - breathing, burning, and decay — would have
filled the air with enough carbon dioxide to destroy life. But com-
pensating devices offset this increase in carbon dioxide.
340
CARBON DIOXIDE
341
All the waters of the earth are continually dissolving carbon di-
oxide, thus removing it i'rom air. In the formation of carbonates and
bicarbonates in nature, great quantities of free carbon dioxide of the
air are used up in the chemical weathering of rocks. Most important
of all, during daylight hours all green plants are absorbing carbon
dioxide, converting it into a form of sugar (fruit sugar) and into
starch, and returning free oxygen to the air. The process that is
responsible for this change, photosynthesis, has been described by
some writers as the most important chemical process in the world,
for if it should cease man and other animals would have no food.
In higher plants, photosynthesis takes place only in parts that con-
tain chlorophyll, a green coloring matter that acts as a catalyst. The
final equation for the formation of starch is:
6CO2 + 5H2O — » 6O2 -f C6H,oO5 (starch)
In Chapter 36 photosynthesis in relation to food-building is dis-
cussed, but at present we are concerned particularly with its effect
upon the composition of air and the way in which it helps to main-
tain the balanced ratio of carbon dioxide to oxygen. This balance,
as it. relates to plant and animal life, is indicated by Fig. 74.
How carbon dioxide is prepared. When an acid is added to a
carbonate or bicarbonate, carbon dioxide is liberated. In the labora-
tory, calcium carbonate, in the form of chips of marble or limestone,
is treated with hydrochloric acid, and the CCX is collected by water
In these tanks, chlorophyll, the catalytic agent in photosynthesis,
is extracted from alfalfa leaves.
American Chlorophyll Division Strong, Cobb and Company, Inc.
342 NEW WORLD OF CHEMISTRY
displacement, as shown in Fig. 75 below. This gas may also be col-
lected by air displacement, for CO2 is heavier than air.
CaCO3 + 2HC1 -» CaCl2 + H2O + CO2 1
Most of the carbon dioxide used commercially is obtained as a
byproduct from coke furnaces, gas wells, and fermentation processes.
The heating of limestone to make quicklime
CaCO3 -» CaO + CO2
also furnishes carbon dioxide for industrial use. Carbon dioxide
for industrial use is separated thoroughly from all impurities before
being liquefied and stored in steel cylinders under pressure. It is
known commercially as carbonic acid gas.
Physical properties of carbon dioxide. Carbon dioxide is a color-
less, odorless gas; it is 1| times as heavy as air, and is soluble in wa-
ter volume for volume (at room temperature and pressure) . At
higher pressures and lower temperatures, it dissolves in water in
much greater volumes.
At ordinary temperatures, a pressure of 52 atmospheres is suffi-
cient to condense carbon dioxide to a colorless liquid. When the
pressure over this colorless liquid is decreased, it evaporates rapidly.
During this process, evaporation of part of the liquid withdraws heat
so rapidly from the remaining liquid that a great amount of the
liquid carbon dioxide is changed to solid carbon dioxide, called
carbon dioxide snow or dry ice.
Chemical properties of carbon dioxide. Carbon dioxide does not
burn. It will support the combustion of only very active substances,
such as sodium. Because it is heavier than air, it can be poured like
a liquid. When a bottle of the gas is poured over a lighted candle,
the flame is extinguished. Why?
The chief chemical property of carbon dioxide is its ability to
combine with water, forming carbonic acid. When carbon dioxide
gas is bubbled through water, part of it combines chemically and
CO2
Fig. 75. Laboratory prepara-
tion of carbon dioxide. What
other gat may be collected
manner?
HCI
Fig. 76. A cubical diagram
of the carbon dioxide mole-
cule.
oxygen
carbon
oxygen
forms carbonic acid; the rest merely dissolves in the water. Carbon
dioxide is, therefore, the anhydride of carbonic acid.
H2O + CO2 -> H2CO3 (carbonic acid)
Electron structure of carbon dioxide. Carbon dioxide is a non-
polar, covalent compound whose structure may be represented as
shown in Fig. 76. Notice that each oxygen atom shares two pairs of
electrons with the carbon atom. The carbon atom shares four pairs of
electrons with the two oxygen atoms. Since the valence of an ele-
ment in a nonpolar compound is equal to the number of shared
pairs of electrons, the valence of oxygen is two and the valence of
carbon is four.
Carbonic acid. Carbonic acid closely resembles sulfurous acid. It
is unstable, decomposing, when warmed, into water and its anhy-
dride, CO2. Thus the reaction for the formation of carbonic acid is
reversible.
H2C03 *± H20 + C02
Because of the highly unstable nature of carbonic acid, it has never
been isolated in the pure state.
Like sulfurous acid, carbonic acid is a weak acid, dissociating only
slightly, forming few hydrogen ions. Carbonated waters containing
dissolved carbon dioxide are so weakly acid that they may be con-
sumed in large quantities without ill effect. Although carbonic acid
turns blue litmus paper pink, its sour taste is scarcely noticeable.
Some of the many uses of carbon dioxide. The chief uses of carbon
dioxide are: (1) in the manufacture of effervescent beverages, (2) as
dry ice for refrigeration, (3) in the leavening of bread, (4) in extin-
guishing fires, (5) in the manufacture of washing and baking soda,
and (6) in the modern synthesis of urea, (NH2) 2CO, a white crystal-
line substance used in the impregnation of wood, in the treatment
of wounds (under the name carbamide) , and in making a valuable
fertilizer, "Uramon":
2NH3 + C02 -> H20 + (NH2)2CO
The normal human adult eliminates about 30 grams of urea as a
waste product daily in liquid excretions.
343
344 NEW WORLD OF CHEMISTRY
Rubber life rafts and life belts are inflated by means of carbon
dioxide. The carbon dioxide so used is stored under pressure in small
steel bulbs not unlike the carbon dioxide cartridges used on siphons
in making charged water.
Another use of carbon dioxide is in making car bo gen, a mixture
of 95 percent oxygen and five percent carbon dioxide, administered to
victims of gas poisoning and pneumonia to induce more rapid res-
piration. Carbon dioxide in the blood appears to stimulate the re-
spiratory nerve centers that control breathing. The heavy breathing
of a runner, for example, is caused by the large quantity of carbon
dioxide produced in his body during violent exercise. Adding more
carbon dioxide to the air inhaled, therefore, causes the more rapid
breathing needed by a patient under treatment.
How carbon dioxide is used in effervescent beverages. Priestley
discovered the pleasant taste of water containing dissolved carbon
dioxide. Before the Royal Society of England he prepared "a glass
of exceedingly pleasant sparkling water which could hardly be dis-
tinguished from Seltzer water," and received the Society's gold medal
for his discovery. This was the first great and unforgettable triumph
of this amateur chemist.
Years later, when he was forced to come to America, Priestley in-
terested Dr. Physick, of Philadelphia, in this beverage. In 1807,
Dr. Physick had a chemist prepare carbonated water with a little
fruit juice for his patients. This was the beginning of the soft drink
industry in America, which uses annually more than 200 million
pounds of carbon dioxide.
Soda water, so named because sodium carbonate was then used
in preparing the carbon dioxide, is now prepared by forcing carbon
dioxide gas into cold water at high pressures. When the pressure is
released, the excess carbon dioxide gas is liberated, causing the bub-
bling, or effervescing. Carbon dioxide is supplied to soda fountains
in liquid form in steel cylinders.
Dry ice. White, solid carbon dioxide has come into wide use for
the refrigeration of foods especially perishable products in transit.
It has three advantages over ice: (1) the temperature of dry ice,
— 78°C., is much lower than that of ice; (2) dry ic$ does not melt
into a liquid, but changes directly into a gas; and (3) in changing
from a solid to a gas, dry ice absorbs three times as much heat ^s ice
does when it melts.
Some dry ice is used in the cooling, and hardening of rivets made
of aluminum alloys, and in shrink-fitting. This is |a process by which
a metal fitting of correct size is expanded by hel$i|kg until it can
Blocks of carbon dioxide dry ice,
direct from the hydraulic presses,
are cut into cubes before being
wrapped for shipment.
•ill Corpnr
be placed over or around the base to which it is to be attached. The
fitting is then shrunk by cooling until it adheres to the base with
great pressure. Dry ice is also used in the low-temperature drying of
many different kinds of biological materials, in the preservation and
shipment of blood plasma, and in the quick-free/ ing- of many dif-
ferent kinds of foods.
Baking powders liberate carbon dioxide lor leavening. Bread
that is not porous is hard, unpalatable, and somewhat indigestible.
Leavening bread makes it porous, light, and more easily digested.
Although bread may be leavened by beating air into the dough, the
method more generally used is to liberate large volumes of carbon
dioxide gas in the dough by chemical action. Carbon dioxide is an
ideal gas for leavening because it is colorless, odorless, nonpoisonous,
and easily and inexpensively prepared.
It is generally produced by baking powders, which are mixtures
of two white powders, one of which is sodium bicarbonate, or baking
soda, and the other a substance such as monocalcium phosphate,
Ca (H,,PO4) .„ cream of tartar, KHC4H4O,., or sodium aluminum sul-
fate, NaAl (SO4) 2. While dry, the two powders do not react, since
most inorganic reactions take place between ions. However, when
water is added, the powders dissolve, dissociate, and an ionic reaction
takes place with the liberation of carbon dioxide gas, causing the
reaction to go to completion. One of these reactions may be repre-
sented as follows:
Baking soda + cream of tartar — > Rochelle salt
NaHCO3 + KHC4O6 -> KNaC,H4O6 + H2O + CO2 \
Sodium bicarbonate + potassium acid — > potassium sodium
tartrate tartrate
Baking powders differ mainly in their speed of action. They also
contain about 15 percent starch or flour as a filler to keep the salts
dry and thus prevent them from reacting before being added to the
dough.
345
346 NEW WORLD OF CHEMISTRY
concentrated H2SO4
Fig. 77. Cross section of a portable
carbon dioxide fire extinguisher. concentrated
solution of NaHCO3
copper tank -»
Some housewives prefer to use their own ingredients in leavening.
For example, in making sour-milk biscuits, which many people pre-
fer to baking-powder biscuits, the housewife uses baking soda and
sour milk, which contains lactic acid.
Carbon dioxide for leavening is produced also by yeast, one of
the oldest of leavening agents. Yeast consists of living plant cells,
which produce zymase. Zymase, which is a mixture of several enzymes,
acts catalytically on the sugar present in the dough and breaks it
down into alcohol and carbon dioxide (see page 548) . What advan-
tage has yeast over commercial baking powders?
How carbon dioxide is used in firefighting. Carbon dioxide gas is
an ideal firefighter. It does not support combustion, is heavier than
air, and can be quickly and cheaply liberated in large volumes.
One kind of common portable fire extinguisher consists of a copper
tank partly filled with a concentrated solution of sodium bicarbonate.
Resting on a shelf inside the top of the tank is a bottle of concen-
trated sulfuric acid covered with a loose lead stopper. When the tank
is inverted, the sulfuric acid pours out of the bottle, and reacts with
the sodium bicarbonate, liberating carbon dioxide gas, which carries
out of the nozzle with it a fine spray of water and some sodium sul-
fate. The equation is:
2NaHCO3 + H2SO4 -> Na2SO4 + 2H2O + 2CO2 1
To put out stubborn fires, such as oil conflagrations which water
cannot extinguish, "Foamite-Firefoam" mixture is used. In this, an
aluminum sulfate solution takes the place of the sulfuric acid in the
ordinary type of fire extinguisher. This aluminum salt combines
with water by a process called hydrolysis, forming gelatinous alumi-
num hydroxide. The carbon dioxide liberated is held fast in large,
tough bubbles by the aluminum hydroxide as well as by a sticky
extract of licorice. Thus, the mixture spreads over the fire a layer
of large bubbles of carbon dioxide. These smother the fire by keep-
ing out the oxygen of the air. The equation is:
A12(SO4)3 + 6NaHCO3 -> 2A1(OH)3 + 6CO2 + 3Na2SO4
CARBON DIOXIDE 347
Fighting fire with liquid carbon dioxide. Another form of fire-
fighting apparatus used very widely depends on liquid carbon diox-
ide under pressure in steel containers. The liquid carbon diox-
ide may be in portable cylinders or in built-in systems (such as in
airplane engines) . When the valve is opened, the liquid carbon di-
oxide vapori/es and rushes very rapidly out of the tube. This sud-
den evaporation of part of the carbon dioxide cools the remaining
liquid to a white solid, and this carbon dioxide snow, played over
the fire, quickly puts it out. The fact that carbon dioxide is able to
penetrate obstructions without damaging equipment makes it one
of the most rapid and efficient of fircfighting substances.
Many ships, including the Queen Mary, are equipped to use car-
bon dioxide for fighting fires in the various storage compartments
throughout the ship. In a recent year, more than 500 million pounds
of liquid carbon dioxide and carbon dioxide gas were produced in
this country.
Salts of carbonic acid. Because carbonic acid contains two replace-
able hydrogen atoms, it is a dibasic acid. As you know, one or both
of these atoms may be replaced by a metal. Hence carbonic acid
forms two series of salts, carbonates and bicarbonates. Thus carbonic
acid reacts with sodium hydroxide, forming sodium carbonate or
sodium bicarbonate, depending upon the conditions.
H2CO3 + 2NaOH -> Na2CO3 (sodium carbonate) + 2H2O
H2CO3 + NaOH -> NaHCO3 (sodium bicarbonate) + H2O
With the exception of sodium carbonate, potassium carbonate, and
ammonium carbonate, all carbonates are insoluble in water. When a
Demonstrating the effectiveness
of a carbon dioxide extinguisher
against a fire in electrical equip-
ment.
348
NEW WORLD OF CHEMISTRY
carbonate is heated, carbon dioxide is liberated and the oxide of
the metal remains.
CaCO3 -» CaO + CO2 f
The carbonates of sodium and calcium are discussed in Chapters 30
and 31.
How to test for a carbonate or bicarbonate. An acid liberates
carbon dioxide gas from either a carbonate or a bicarbonate. If this
gas is then passed into limewater, Ca (OH) 2, a white precipitate, cal-
cium carbonate, CaCO3, forms.
Ca(OH)2 + C02 -> CaC03 1 + H2O
On bubbling more carbon dioxide gas through the mixture, the
white precipitate disappears as a result of the solubility of calcium
bicarbonate which is formed.
CaCOg + H2O + CO2 -> Ca(HCO3)2 (calcium bicarbonate)
YOU WILL ENJOY READING
Armstrong, George B. "Dry Ice." Chemistry, Feb. 1951,
pp. 24-31. An excellent, illustrated article on the history,
properties, manufacture of and uses of solid carbon dioxide.
U.S. Department of Agriculture, 1945. Safe Use and Storage
of Gasoline and Kerosene. Farmers' Bulletin No. 1678. Supt. of
Documents, Washington, D.C. 10^. Some excellent advice on
putting out gasoline and kerosene fires.
USING WHAT YOU HAVE LEARNED
1.
2.
3.
4.
write
by.,
and .
5.
cycle.
6.
Group A
What are the chief sources of CO2?
How and by whom was CO2 discovered?
In what way does life depend upon CO2?
Copy and complete the following statements. Do not
in this book. CO2 is being added to the air constantly
. , . . . , and .... CO2 is removed from the air by . . . , . . . ,
Make a diagram representing the carbon dioxide-oxygen
By what three commercial methods is CO2 produced?
CARBON DIOXIDE 349
7. Write the equation for one commercial method of pro-
ducing CO2.
8. Make a labeled diagram for the laboratory preparation
of CO2.
9. (a) Write the chemical equation for the laboratory
preparation of CO2. (b) Why does this reaction go to com-
pletion?
10. (a) Can carbon dioxide be collected by air displace-
ment? (b) Explain.
11. What are the physical properties of CO2?
12. What are three chemical properties of CO2?
13. (a) Write the reversible reaction of carbon dioxide and
water, (b) How could you make large quantities of CO2 com-
bine with H2O? (c) The behavior of CO2 with water places it
in what class of compounds?
14. What other acid closely resembles H2CO3?
15. (a) What are six uses of CO2? (b) Opposite each use
state the property or properties on which that use depends.
16. Describe briefly the beginning of the soft-drink industry
in America.
17. (a) How is dry ice made? (b) Dry ice has what advan-
tages over ice? t
nr^"
18. What is the general composition of baking powder?
19. What is the function of each ingredient in one type of
baking powder?
20. Write a balanced equation representing the liberation
of carbon dioxide from one type of baking powder.
21. (a) Describe the chemical action of some other type of
leavening agent, (b) What advantage has it over commercial
baking powders?
22. What are four reasons that CO2 is an ideal leavening
agent?
23. A manufacturer makes enough baking powder to pro-
duce 15 tons of CO2. What weight of NaHCO3 does he use?
24. A sample of baking powder liberated 200 ml. of CO2.
Considering that the reaction went to completion, what weight
of NaHCO3 was used if the baking powder was of the cream-
of- tartar type?
35O NEW WORLD OF CHEMISTRY
I
25. What volume of air would be needed to burn completely
enough coke to make 10,000 cu. ft. of CO2?
26. Why do baking powder reactions go to completion?
27. Describe with diagram and equation the principle of
one type of portable fire extinguisher.
28. (a) H2O is not a good agent for putting out a fire of
burning oil. Why? (b) Why is CO2 from the ordinary type of
fire extinguisher ineffective? (c) What principle is used in the
"Foamite-Firefoam" method?
29. (a) H2CO3 has two series of salts. What are they?
(b) Write two equations to illustrate.
30. What is the action of heat on carbonates?
31. (a) Describe a complete test for a carbonate, (b) Write
the equation for this chemical test.
32. When sour milk is used in cooking, NaHCO3 is used
instead of, or in addition to, baking powder. Explain.
33. How can CO2 be removed from a sample of air?
34. Can pure GO2 alone be used in fighting fire? Explain.
Group B
35. (a) Why does CO2 collect in wells and caves? (b) What
is its source?
36. Does the effervescence of a solution when an acid is
added prove the liberation of CO2? Explain.
37. 250 ml. of a gas weigh 0.49 g. This gas contains 27.3
percent carbon and 72.7 percent oxygen. Find the formula of
this compound.
38. Most chemical reactions take place between ions. How-
ever, CaCO3 is insoluble in water. Explain the liberation of
CO2 when HC1 is added to solid CaCO3.
39. Explain why a knowledge of chemical arithmetic is es-
sential in preparing baking powders.
40. SO2 is a reducing agent but CO2 is not. Explain.
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Make a study of the cost per ounce of a half-dozen widely
advertised baking powders. Nearly all baking powders liberate
CARBON DIOXIDE 351
the same amount of CO2 for equal weights of baking powder.
Is the wide range of prices for such baking powders warranted?
Give reasons for your answers. Why do such wide ranges in
prices exist?
2. Question your mother, or other members of your family,
and friends about their preferences in leavening agents. List
the data you obtain and give your own reactions to such pref-
erences.
3. Visit a soft-drink bottling plant in your neighborhood,
and report on the method used in charging the soft drinks
with CO2.
4. After consulting your teacher of biology on sources of
information, prepare a two- or three-page report on micro-
organisms, other than yeast, that are used by man in the prep-
aration of his food supply. Arrange a class discussion on this
topic.
5. Construct a glass model of an acid-sodium bicarbonate
type of portable fire extinguisher. Demonstrate its action be-
fore the class.
CARBON MONOXIDE:
OTHER GASEOUS FUELS
. . . The introduction of gas light-
ing (1810) met with much opposi-
tion, as one can understand, on the
part of the tallow chandlers and sell-
ers of whale oil — at that time used
as an illuminant; but after a time
the new method of lighting was wel-
comed. Alexander Findlay, 1930
Priestley discovers an oxide of carbon that burns. Almost half a
century after the discovery of carbon dioxide, another quite differ-
ent oxide of carbon was found, carbon monoxide, CO. Not until the
last year of the eighteenth century was carbon monoxide known to
be a new compound of carbon.
Priestley again was the first to recognize clearly a new gaseous
compound in the flickering blue flame which played over his fur-
nace fire. This valuable discovery was made in America, where this
dissenting English minister had fled to escape political and religious
tyranny.
How carbon monoxide is formed. Priestley's observation indicated
that when coal or other carbonaceous fuel is burned in a limited sup-
ply of air, carbon monoxide, rather than carbon dioxide, is formed.
2C + O2 -> 2CO
Carbon monoxide is formed also in a house furnace by the reduc-
tion of carbon dioxide gas as it passes over glowing coal, which, as
you know, is an excellent reducing agent.
CO2 + C -» SCO
CARBON MONOXIDE 353
Normally, carbon monoxide burns at the surface of a coal fire,
giving off heat and changing to carbon dioxide, but if the supply of
air is insufficient, or if the flame is chilled, some of the carbon mon-
oxide may escape without burning.
How carbon monoxide is prepared in the laboratory. Formic acid
is a colorless liquid with a characteristic odor, that is, a distinct odor
that can be used to identify it. Its formula is HCOOH (only the
last H is a replaceable hydrogen atom) . When warm concentrated
sulfuric acid, a powerful dehydrating agent, is added to formic acid,
it extracts the HOH of the acid, leaving pure CO.
HCOOH -> HOH + CO t
Physical properties of carbon monoxide. Carbon monoxide is a
colorless, tasteless, almost odorless gas, slightly lighter than air, and
slightly soluble in water. Its odor is very difficult to detect. It can be
liquefied, but only under high pressures and low temperatures.
Chemical properties of carbon monoxide. Carbon monoxide burns
in air with a pale blue flame, forming carbon dioxide, which, because
it is unable to unite with any more oxygen, cannot burn.
2CO + 02 -» 2C02
Because of carbon monoxide's great attraction for oxygen, it is
an excellent reducing agent, and therefore it is used in the extraction
of iron from its oxide. The carbon monoxide is formed by the reac-
tion of coke (carbon) and CO2 in the upper part of a blast furnace.
Fe2O3 + SCO -> 2Fe + 3CO2 T
Carbon monoxide combines with certain metals, forming a series
of compounds called carbonyls. For example, when nickel is heated
in the presence of carbon monoxide to a temperature of about
40°C., nickel carbonyl, a poisonous liquid, is formed.
Ni + 4CO -> Ni(CO)4 (nickel carbonyl)
Upon further heating, carbonyls liberate CO and give up their met-
als. Nickel carbonyl is used in the extraction of nickel from its ores.
Carbon monoxide in the presence of suitable catalysts also com-
bines directly with hydrogen, forming wood alcohol (see page 548) ,
and with chlorine, forming phosgene.
CO + 2H2 -» CH3OH (wood alcohol)
CO + C12 — > COC12 (phosgene)
354 NEW WORLD OF CHEMISTRY
Some uses of carbon monoxide. Carbon monoxide is an excellent
fuel, burning with intense heat and leaving no solid residue. It is
most often used in conjunction with other gases as a mixed fuel, as
in water gas, for example, which is a mixture of carbon monoxide
and hydrogen.
Carbon monoxide is used widely in the metallurgy (extracting a
substance from its ores) of iron and nickel.
Deadly carbon monoxide, a cumulative poison. Carbon monoxide
is the most widespread poison connected with human life and activ-
ity. It acts painlessly. Furthermore, it gives no warning of danger,
because it is almost odorless, and its victim slowly passes into a state
of unconsciousness. Its physiological action is caused by the strong
affinity of carbon monoxide for hemoglobin, an attraction that is 300
times greater than the attraction of oxygen for hemoglobin.
Hemoglobin is present in the red blood corpuscles and its function
is to unite chemically with oxygen and carry it to all parts of the body
by the circulation of the blood. By forming a stable compound with
the hemoglobin in blood, carbon monoxide prevents oxygen from
reaching millions of cells in the body. The victim is thus killed by
suffocation — a lack of oxygen.
The formation of carbon-monoxide-hemoglobin takes place gradu-
ally, and this substance remains in the blood. When the percentage
of this stable compound has reached 40 percent, the victim collapses;
when it reaches 80 percent, death occurs. Such a poison, which col-
lects slowly in the body and over a long period of time until it reaches
a dangerous concentration, is said to be a cumulative poison.
The treatment for carbon monoxide poisoning consists of induc-
ing respiration by the Schafer prone-pressure or the more recently
adopted Holgar-Nielsen "back-pressure, arm-lift" method in the open
air, while carbogen is administered. No alcohol should be given. As
you know, the carbon dioxide makes the patient breathe more deeply.
The injection into the blood stream of methylene blue, a dye, as an
antidote for carbon monoxide poisoning has been reported to be
effective in certain cases.
Where the danger of carbon monoxide poisoning lurks. The ex-
haust gas of automobiles, in which carbon monoxide exists in vary-
ing percentages up to about 12 percent, is one of the most common
sources of carbon monoxide. Hundreds of persons are killed by this
gas every year, usually in some such ways as these.
1) A motorist drives his car into his garage. It is cold, and he shuts
the doors of his garage to keep warm. He keeps the motor running
while he gets under the machine to make some repairs. Before he is
CARBON MONOXIDE
355
aware of it, carbon monoxide from the exhaust pipe of the automo-
bile causes unconsciousness. Death soon follows.
2) On a very cold morning a driver may start the engine before
opening the garage door, or he may sit in a closed car with the en-
gine running while waiting for a friend. Carbon monoxide gas from
the exhaust finds its way into the confined space and takes its toll.
The carbon monoxide comes from the incomplete combustion of
the gasoline. Although gasoline is a mixture of hydrocarbons, we may
represent its composition by the formula C7H16. During incomplete
combustion, the following reaction is one that may take place:
C7H]6 + 8O2 -» 6CO + 8H2O + CO2
The most serious industrial cases of carbon monoxide poisoning
occur from exposure to gases, such as producer and blast furnace
gases used in many manufacturing plants (see pages 359 and 407) .
Most poisonings from carbon monoxide outside of industrial and
auto cases are caused by the escape of gas from kitchen ranges and
stoves, gas refrigerators and other gas appliances because of leaks
and the accidental extinguishing of the burning gases. A furnace, im-
properly adjusted, may give off quantities of this gas and become a
menace to life. As shown below, every coal stove and furnace should
have an adequate circulation of air, otherwise carbon monoxide may
be formed and finally reach dangerous concentrations.
Firing for complete combustion. In firing a coal-burning furnace,
the object is to obtain as much heat from the coal as possible. This
can be done only when combustion of the coal is complete, or nearly
so, with a minimum of intermediate products, including carbon
monoxide and free carbon, or soot, remaining unburned. Incom-
plete combustion is dangerous because of the dangers from the car-
bon monoxide produced. It is also wastefully inefficient and robs
chimney
air
Fig. 78. Stove in operation.
When damper A is open, but
dampers B and C are not, in-
complete combustion results.
When 0 and C only are open,
the fire burns slowly.
2C02'
air
of
burning CO
reduction of CO2
combustion
of coal
grate
356 NEW WORLD OF CHEMISTRY
the consumer of many Btu per dollar of fuel. In addition, the great
quantities of smoke and soot produced are a menace to health and
greatly increase expenditures for laundry and dry cleaning.
Complete combustion of coal depends upon adequate quantities of
air and upon proper firing methods, that is, proper placement of coal
in the firebox. In adding coal to a burning fire, push the glowing coals
from the front of the firebox to the rear, and place the additional coal
in the space from which the coals were pushed. In this way, the fire is
not smothered and great clouds of smoke do not result. The glowing
coals at the rear of the firebox heat the added coal, combustible va-
pors are formed and oxidized relatively completely in the region sep-
arating the new coal from the glowing coals.
If the new coal is spread completely over the glowing coals the
fire is partially smothered. Combustible vapors are formed, to be
sure, but for a long period, the temperature above the added fuel is
lower than the kindling temperature of the vapors, and very small
amounts of heat are produced.
Proper regulation of the air supply is a most important factor in
securing complete, or nearly complete, combustion. If the supply of
air is insufficient, incomplete combustion results, and carbon mon-
oxide and free carbon, or soot, are produced in large quantities. If
too much air is available, combustion is relatively complete but very
rapid, and results in uneven heating. If far too much air is available,
much of the heat produced is carried out the chimney without trans-
ferring its heat to the water, steam, or air in the heating system.
Finding the damper settings that result in relatively complete com-
bustion is not difficult, but it does require some patience and much
observation, for under varying outside conditions, damper settings
vary also. If you can adjust the dampers so that the flame produced
in the firebox has a minimum of yellow flame, the conditions neces-
sary for nearly complete combustion probably are present.
Coal fires Should not be poked or shaken as often as most persons
poke and shake them. Shaking of ashes should occur only when the
accumulation of ashes becomes so deep that the air supply may be im-
peded. Ashes insulate the metal grates of the firebox from the in-
tense heat of the fire, and if they are all removed, the grates may be
damaged as a result of receiving more heat than they were designed
to receive. Poking a fire frequently spreads the burning coals through-
out the firebox and results in too rapid combustion^
The simple principles just discussed will, if followed, do much
to increase the heat that consumers get from the c6al they buy; smoke
nuisances will be reduced, and fuel bills will be ctil itiaterially.
CARBON MONOXIDE
357
How carbon monoxide can be removed from air. For a long time
no simple method of removing noxious carbon monoxide from air
was known. Activated charcoal, which adsorbs various gases, such as
ammonia, acetone, benzene, and chlorine, will not remove carbon
monoxide from air.
A catalyst called "Hopcalite," which consists of a mixture of me-
tallic oxides, causes carbon monoxide to change to carbon dioxide
at ordinary temperatures and pressures. In the presence of "Hop-
calite," 1 molecule of carbon monoxide combines with another atom
of oxygen from the air. A rescue-breathing apparatus containing a
"Hopcalite" canister is now used by persons who find it necessary
to enter regions where concentration of carbon monoxide is high.
How can carbon monoxide be detected? Simple color-detector
tubes have been devised for determining minute amounts of carbon
monoxide in air. One such preparation called hoolamite contains
specially prepared iodine pentoxide, L()r>, which oxidizes carbon
monoxide to CO2.
5CO + I2O& -» 5CO2 + I2
The iodine fumes liberated cause a change in color that is directly
proportional to the amount of carbon monoxide present. This
amount can be determined by comparison with a standard color
scale. The color fades, and the tubes may be used again. Before this
method was developed, canaries were often used as detectors because
these birds are very sensitive to minute amounts of this poison and
show its effects before man does.
Official U.S. Navy photograph
Navy firefighters ready for
action. Their rescue-breathing
apparatus will protect them
from carbon monoxide and
other dangerous fumes.
These recording devices keep
a constant check on the con-
dition of the air in the Lincoln
Tunnel. The 8000-foot vehi-
cular tunnel connects the
New York and New Jersey
banks of the Hudson River.
The Port of New York Authority
In some tunnels, such as the Holland and Lincoln Tunnels at
New York City, machines have been installed which record on a
time chart the amount of carbon monoxide present in the air. A
warning bell is caused to ring when the carbon monoxide reaches
a concentration of four parts in 10,000. At this point, also, dampers
and fans begin to operate automatically to change the air.
What are the more important gaseous fuels? The fuel gases
burned each year in this country are worth more than two billion
dollars. The total volume of gas produced is about eight trillion
cubic feet. Besides pure carbon monoxide, which is seldom used
alone, the most important gaseous fuels are water (or synthesis) gas,
producer gas, coal gas, natural gas, acetylene gas, and hydrogen. Hy-
drogen has already been discussed; a brief discussion of the others
follows.
A gas from steam and carbon. Water gas is manufactured by pass-
ing water in the form of steam over glowing coke or hard coal. The
coke is first burned in a draft of air until it is red-hot. The air is
then shut off and the steam turned on. The temperature of the coke
falls gradually because the reaction is endothermic, and when it
reaches about 1000°C., the steam is cut off and the air supply re-
newed. During the process the carbon, which is an ideal reducing
agent, combines with the oxygen of the water, leaving hydrogen.
H2O + C -> CO + H2
water gas
At a temperature of about 1000°C., the most gas is produced.
Above this temperature, the carbon monoxide formed may react
with the steam and change to carbon dioxide.
H2O + CO -> CO2 + H2
358
CARBON MONOXIDE 359
Water gas is thus a mixture of carbon monoxide and hydrogen.
It is used, either alone or mixed with coal gas, for domestic heating
purposes. Since both carbon monoxide and hydrogen burn with al-
most colorless flames, a mixture of the two cannot be used for illumi-
nating purposes unless it is made luminous by injecting into it gas-
eous hydrocarbons from petroleum, which, on burning, give a yellow
flame. This process of adding hydrocarbon vapors to water gas is
called enriching the gas. When water gas burns, it forms water va-
por and carbon dioxide.
CO + H2 + O2 -» H2O + CO2 T
water gas
When used in the synthesis of gasoline and other chemicals it is
called synthesis gas.
The gasification of coal. In areas where natural gas is not available
as a fuel for factories and homes, producer gas is most generally
used. It is produced by burning low-grade coal in a furnace with a
limited supply of air. The chief product formed during the incom-
plete combustion of the coal is carbon monoxide. The gas issuing
from this furnace is mixed with large quantities of nitrogen, which
is too inactive to unite with the coal. Producer gas, then, is chiefly
a mixture of about 60 percent nitrogen and 30 percent carbon mon-
oxide. It contains also about ten percent hydrogen.
2C + air (02 + N2) -» 2CO + N2
Producer gas is a much poorer fuel than water gas because it con-
tains such a large amount of nitrogen, an incombustible gas. Its
manufacture constitutes the most efficient means of converting low-
grade coal into power.
A gas similar in composition to producer gas was first proposed by
Mendeleyeff and later produced in 1933 in the Soviet Union by
burning coal underground. Instead of mining the coal in the usual
way, a coal seam was sealed off, set on fire, and the gases produced
were brought to the surface through pipes. Later, pure oxygen was
mixed with the blast of air sent into the mine, and the gas that came
out of the mine was richer in composition than producer gas.
This gasification of coal underground (gas mining) , held feasible
by the U.S. Bureau of Mines after several years of experimentation,
may turn out to be a revolutionary development. By this method we
may be able to recover valuable coal in mines abandoned because of
the low-grade nature of the coal or because of the thinness of the coal
360
Fig. 79. Underground gasiflcc
tion of coal.
air
producer
gas
gasified
area
vein. In addition, this method makes coal working safer and health-
ier, a consideration which is always desirable. Another advantage is
pipeline distribution of this fuel.
Natural gas, the Cinderella fuel. More than 80 percent of the
gaseous fuel consumed each year in the United States is natural gas.
In various sections of the world, especially where coal and petroleum
deposits are found, natural gas issues from the earth when porous
rocks saturated with it are tapped. The gas may flow out under pres-
sure, or pumping may be required.
Many gas wells do not yield oil, but an oil well almost always
produces both gas and oil. The gas from wells that do not yield oil is
usually very rich in methane, CH4, some wells yielding as high as
95 percent methane.
In the early days of the petroleum industry, little or no use was
found for natural gas, and most of the wells were ignited and allowed
to burn for years. When the waste of natural resources involved in
such practices was realized, controls were established and the burn-
ing of natural gas greatly diminished.
Methane, also called marsh gas, is a colorless, practically odorless,
insoluble gas which burns with an almost colorless flame, forming
CO2 and water vapor.
CH4 + 2O2 -> CO2 + 2H2O
It has a calorific value about twice that of manufactured gas which
it has largely replaced. It is about half as heavy as air. Natural gas is
purified, some high quality casing-head gasoline being obtained in
CARBON MONOXIDE
361
this process. H2S is also removed before it is sent into pipelines. It is
also stripped of its heavier hydrocarbons such as butane and propane
before being piped many hundreds of miles to supply factories, farms,
and homes with light, heat, and power. For example, there is an
1840-mile pipeline (the Big Inch) from Texas to New York.
Probably the origin of our great supply of natural gas is the same
as that of petroleum, which most scientists believe to be the result of
the incomplete decomposition of vegetable or animal matter, either
with or without bacterial action. The formation ,of petroleum does
not require millions of years as was formerly believed. This is demon-
strated by the formation at the present time of petroleum in off-shore
marine sediments. Some scientists believe, however, that natural gas
originates from the interaction of metallic carbides and water, just as
acetylene is formed by the reaction between calcium carbide and
water.
CaC2 + 2H20 -» Ca(OH)2 + C^ t (acetylene)
A gas from a gray solid. The gaseous fuel called acetylene, C2H2,
is colorless and odorless when pure, very slightly soluble in water,
and somewhat toxic. Acetylene has a tendency to explode when lique-
fied. The gas is therefore not liquefied, but is forced at low pressure
into a solvent called acetone, a colorless liquid obtained from the
destructive distillation of wood, and mixed with some inert porous
material, such as wood charcoal or asbestos fiber. It is sold in port-
able steel cylinders and is used widely in oxyacetylene torches. When
the valve of one of these cylinders is opened, the pressure is reduced,
and some of the gas escapes from solution.
For emergency use, and in lighthouses and isolated districts where
electric lighting and illuminating gas are not available, special acety-
lene generators have been constructed. These special generators al-
low water to come in contact with calcium carbide at regulated rates
of speed, so that the gas may be liberated as needed.
362
NEW WORLD OF CHEMISTRY
CaC2
acety
Fig. 80. (left) Carbide-to-water acetylene generator for large installations, (center)
Water-to-carbide generator, (right) The acetylene burner.
The burner used with acetylene gas must be specially constructed
to permit the access o£ a large amount of air to the burning gas.
Because acetylene is rich in carbon, it would otherwise burn incom-
pletely and produce a smoky flame.
2C2H2 + 5O2 —> 4CO2 + 2H2O (complete combustion)
2Q>H2 + O2 — > 4C -h 2H2O (incomplete combustion)
Alcohol, acetone, vinegar, chloroform, plastics, and synthetic fibers,
rubber, and gasoline have been built up chemically from acetylene
gas (see Chapters 34 and 35) .
A gas from the destructive distillation of coal. About the time
that Priestley was studying carbon monoxide, coal gas was intro-
duced as an illuminant. William Murdock, a Scottish workman em-
ployed by James Watt, developer of the steam engine, carried out
experiments that led to lighting part of the Boulton and Watt fac-
tory in Birmingham, England, with gas in 1798. A few years later,
gaslighting was introduced in the United States. Much opposition
was raised against it, but the advance of science could not be stopped
for long. Manufacturers of candles and whale oil, fought against il-
luminating gas, even as other manufacturers have since struggled un-
successfully against other innovations.
Coal gas is obtained from coal by destructive distillation. Fig. 81
shows the method of manufacture and purification in a byproduct
coke oven. Coal is heated in a closed oven. The vapors formed are
first passed into the hydraulic main, where some of the vapors of
impure gas
Fig. 81. Steps in the production of coal gas.
gas
holder
CARBON MONOXIDE 363
coal tar and ammonia are condensed. The remaining vapors then
enter the condensers, where the rest of the ammonia is absorbed and
the coal tar condensed. The purifier, containing iron oxide or lime,
removes hydrogen sulfide and other sulfur compounds. The purified
coal gas then enters the gas holder. Coal gas contains about ten per-
cent carbon monoxide, 40 percent hydrogen, and 40 percent methane.
It contains also about ten percent nitrogen.
The destructive distillation of one ton of coal yields approxi-
mately 10,000 cubic feet of coal gas, 20 gallons of ammonia water,
and 120 pounds of coal tar. About 1400 pounds of coke remain in
the retort. A tarry matter called pitch is left also, which is used as a
binder in road construction.
What is an explosive mixture? All the gases discussed in this
chapter burn quietly in air. However, when they are mixed with air
in the right proportions to secure nearly complete combination and
then ignited, very rapid oxidation takes place, suddenly producing
extremely large volumes of gases. The high temperature of the re-
action helps to account for the large volume of gases formed, since
gases expand as their temperature is raised. High temperatures also
result in high pressures. If the explosive mixture is confined, as, for
example, in a mine or factory, great destruction takes place.
Gases and vapors differ from one another in the range of composi-
tion of their explosive mixtures. Thus coal gas will explode when
anywhere from six percent to 29 percent of it is mixed with air. Air
containing less than six percent or more than 29 percent of coal gas
will not explode.
The character of the explosive mixture is of prime importance in
the working of internal-combustion engines. For example, the degree
of smoothness and of power in the running of an automobile engine
depends upon getting the right mixture of air and gaseous fuel
admitted to the cylinders and ignited at the proper instant by an
electric spark. The closer the proportions of the mixture to those
necessary for securing complete combustion, the greater the power
produced and the less the waste of fuel.
Because of the importance attached to securing this proper mix-
ture, the carburetor, in which the mixing is done, is often called the
heart of the internal-combustion engine. This needs adjustment from
time to time to insure highest efficiency (see Fig. 82) .
Gasoline vapor mixed with air may form a very explosive mixture.
Many serious accidents have resulted from the careless use of gaso-
line in dry cleaning at home. Such cleaning operations should be
done in the open and away from any source of ignition. In using
364
NEW WORLD OF CHEMISTRY
mixture of
gas and air
carburetor
cylinder
piston
spark plug
exhaust
Fig. 82. Carburetor and cyl-
inder assembly. The gaso-
line-air mixture is ignited in
*he cylinder by the spark
P|ua- The «neray of tne ex-
plosion is transformed into
motion by the piston.
gasoline for dry cleaning, vigorous or continued rubbing of the fab-
ric should be avoided lest the friction ignite the gasoline.
Safety measures against mine explosions. For centuries the igni-
tion of explosive mixtures, especially in coal mines, has caused seri-
ous loss of life. In 1556, Agricola («-grik'6-la) published a book on
mining, one section of which dealt with "the ailments and accidents
of miners and the methods by which they can guard against these."
But life was cheap in those days and little was done to protect min-
ers against explosions. At the end of the eighteenth century, there
came an emphasis on the rights of man and with it a new humani-
tarianism. An interest in occupational accidents and diseases was
aroused. Gases in mines and mine ventilation were studied.
Davy, who discovered sodium and potassium, and made other im-
portant contributions to chemistry, was among those interested in
the plight of the miners in England. He devised a simple safety lamp
to prevent mine explosions. It is based upon two principles: (1) An
explosive mixture does not undergo chemical change until its kin-
dling temperature is reached. (2) Metal surfaces spread heat rapidly.
In the Davy lamp, the flame is surrounded by a wire gauze, which
distributes the heat produced by the flame over a wide area, and thus
prevents the explosive mixture outside the lamp from reaching its
kindling temperature. A lighted match brought over a wire gauze,
as shown in the illustration below, will set fire to the gas above the
gauze, but the gas below will not catch fire. Why? This is an illustra-
tion of the principle of the Davy safety lamp.
In commercial mining, the Davy lamp has been replaced by a
battery-operated electric lamp.
gas flame
wire gauze
gas below its
kindling temperature
Fig. 83. Demonstration of the
principle of the Davy lamp.
Why does the gas below the
wire gauze not catch fire?
In an underground coal
mine, the air is tested for
the presence of methane
immediately after coal is
blasted from the seam.
The flame of the safety
lamp turns blue when the
gas is present.
National Coal Association
YOU WILL ENJOY READING
Faraday, Michael. The Chemical History of a Candle.
"Kings' Treasuries of Literature" Series, E. P. Button 8c Co.,
New York, 1920. This book consists of six lectures delivered
by Faraday before young boys and girls at the Royal Institution
of London in 18(>()-1861. Lecture I covers (lames — sources and
structure. The book is a classic, and though about 100 years
old may still be read with pleasure and profit.
Manchester, Harland. New World of Machines, pp. 174-
189. Random House, New York, 1945. "Power for Tomorrow"
is a fine chapter in this carefully written book.
Oettingen, W. F. von. Carbon Monoxide: Its Hazards and
the Mechanism of Its Action. Public Health Bull. No. 290,
1944. Supt. of Documents, Washington, D.C. 35^.
USING WHAT YOU HAVE LEARNED
Group A
1. Where and by whom was pure GO first studied?
2. (a) How is CO formed in a furnace? (b) Write an equa-
tion to illustrate, (c) Write the equation for the oxidation of
CO.
3. (a) How is CO usually prepared for laboratory use?
(b) What is the function of the H.,SO4 used?
4. Compare the physical properties of N2 and CO.
5. (a) What are two chemical properties of CO in addition
to its combustibility? (b) Write equations to illustrate each.
6. Write balanced equations for the following. Name the
products, (a) Action of chlorine and carbon monoxide,
(b) action of hydrogen and carbon monoxide, (c) action ol
carbon monoxide on heated nickel, and (d) formation of
phosgene.
365
366 NEW WORLD OF CHEMISTRY
, , | , t T
1
7. Discuss four uses of CO.
8. (a) Explain how CO acts on the hemoglobin of the
blood of a person breathing it. (b) What first-aid treatment
should be given a person who has been overcome by CO?
(c) What is the function of the small amount of CO2 ad-
ministered?
9. Poisoning by CO often occurs in closed garages,
(a) Why? (b) What precautions should be taken to prevent
this danger?
10. (a) What gases are generally adsorbed by the C used in
gas masks? (b) What substance is used in CO detectors?
11. (a) What are two important factors in firing with coal
for complete combustion? (b) What is the most efficient
method of firing? (c) Why?
12. In firing with coal, what are the results of (a) inade-
quate air supply? (b) too much air? (c) far too much air?
13. How can you tell when the conditions necessary for
complete, or nearly complete, combustion are likely to be
present?
14. (a) What do large quantities of black smoke issuing
from a chimney indicate? (b) Why?
15. (a) Is black smoke undesirable? (b) Why?
16. (a) What are two pure gaseous fuels? (b) What are
three that are mixtures of gases? (c) Which of (b) contain
large quantities of CO?
17. (a) Write an equation for the manufacture of water
gas. (b) What is "synthesis gas"?
18. What volume of steam is used in making 1500 cu. ft. of
water gas?
19. What weight of coal, containing 85 percent C, is used
in making 2000 cu. ft. of water gas? (One , ounce-molecular
weight of a gaseous substance occupies 22.4 cu. ft.)
20. Water gas contains 60 percent CO and 40 percent H2.
What volume of air is necessary for tlie complete combustion
of 200 cu. ft. of this gas?
21. (a) What is meant by enriching water gas? (b) Why
is it done?
CARBON MONOXIDE 367
22. Producer gas contains about 60 percent noncombustible
N2, yet it is used in great quantities as a fuel for gas engines,
(a) How does this N2 affect its fuel value? (b) What process
uses producer gas? Why?
23. Explain the gasification of coal underground.
24. (a) On the basis of present evidence, what do we think
was the source of natural gas? (b) How does its present con-
sumption rank with that of other gaseous fuels? (c) What is
its chief constituent?
25. (a) What are the properties of methane? (b) What
other name has it? (c) Why?
26. Make a table showing the source, composition, and chief
use of four gaseous fuels.
27. Describe briefly the stages in the manufacture of puri-
fied coal gas.
28. (a) What is destructive distillation? (b) Name the chief
products of destructive distillation of bituminous coal?
29. (a) How is C2H2 made? (b) What are its principal
properties? (c) Why must a C2H2 burner be so constructed
that it allows the access of large quantities of air?
30. How many cubic feet of air are used in burning com-
pletely 2500 cu. ft. of C2H2?
31. How would you identify CO2, H2, CO, N2, and NO?
32. (a) What happens when an explosive mixture is ignited?
(b) Why are explosive mixtures dangerous? (c) What is the
most important factor to consider in the mixing of gases for
explosive effect?
33. (a) How is the explosive mixture in an automobile
regulated? (b) Why?
34- Describe a situation at home in which a dangerous ex-
plosive mixture of gases might be formed.
35. What might happen if the ventilating system in an
underwater tunnel such as the Lincoln Tunnel in New York
City suddenly got out of order?
Group B
36. Why does CO burn, whereas CO2 does not burn?
37. In making water gas, what would happen if the tempera-
ture used were too high?
38. Why has the byproduct furnace replaced the beehive
coke oven?
368 NEW WORLD OF CHEMISTRY
39. Why is a blue flame seen when fresh coal is added to
burning fuel?
40. A gas sample from a sealed fire area in a mine shows
CO, four percent; O2, ten percent; CH4, seven percent; and N2,
79 percent. Is the fire bla/ing, or is a methane explosion pos-
sible? Explain.
41. Suppose a sample of gasoline is half hexane, C6H14, and
half heptane, C7H16. How many cubic feet of air are necessary
for the complete combustion of 20 cu. ft. of this gasoline
vapor?
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. If you live near an oil or natural-gas field, visit the gas
wells and write a report on the gas waste, if any. Do you think
Federal control of such waste desirable? Explain.
2. Davy refused to take out a patent on his miner's lamp,
saying, "No, my good friend, I never thought of such a thing;
my sole object was to serve the cause of humanity, and it I have
succeeded I am amply rewarded in the gratifying reflection of
having done so." Write a report expressing your views on this
incident. Cite similar instances.
3. Study the recently introduced Holgar-Nielsen "back-pres-
sure, arm-lift" method and compare it with the Schater prone-
pressure method.
4. Using a test tube, a one-hole rubber stopper, a short de-
livery tube and some dry sawdust or wooden splint, show that
a combustible gas can be obtained Irom the wood similar to
coal gas from coal. Consult your laboratory workbook.
25
METALS
AND THEIR CHEMICAL ACTIVITY
. . . Potassium and sodium are the
names by which I have ventured to
call the two new substances . . .
. They agree with the metals in opaci-
ty, lustre, malleability, conducting
powers as to heat and electricity, and
in their qualities of chemical combi-
nation. Their low specific gravity
does not appear a sufficient reason
for making them a new class.
Sir Humphry Davy, 1807
Elements may be classified as metals or nonmetals. Thus far we
have discussed a number of elements including oxygen, nitrogen,
chlorine, bromine, iodine, fluorine, sulfur, and phosphorus. Each
of these elements has more than one-half the number of electrons in
its outer electron ring that are necessary to complete this ring. Hence,
each is a borrower of electrons, has a negative valence, and has an ox-
ide (or oxides) that is an anhydride of an acid. Each of these ele-
ments is called a nonmetal.
But the properties of these elements are not characteristic of all
the elements. Certain elements have less than half the number of
electrons in their outer electron rings necessary to complete these
rings. Hence, each of these elements is a lender of electrons, has a
positive valence, and possesses an oxide (or oxides) that is an an-
hydride of a base. Each of these elements is called a metal. Of the
100 chemical elements, 78 are metals.
A comparison of sulfur and magnesium will help to make clear the
differences between a metal and a nonmetal.
This classification of elements into metals and nonmetals is very
old, but it still continues to be of service. We must remember, how-
ever, that certain elements behave either as a metal or as a nonmetal.
369
Magnesium
a typical
metal
Sulfur
a typical
nonmetal
a) 6 electrons in outside ring a) 2 electrons in outside ring
b) borrower of 2 electrons b) lender of 2 electrons
c) valence is — 2 c) valence is +2
Fig. 84. Structure of the atoms of a typical metal and nonmetal.
How metals occur in nature. Metals occur either free (uncom-
bined) , or chemically combined in compounds. Gold, silver, copper,
and platinum are examples of metals that occur in the free state.
The oxide of iron, the fluoride of aluminum, the chloride of sodium,
the bromide of magnesium, the sulfate of barium, the carbonate of
lead, and the phosphate of calcium are examples of compounds of
metals that are found in the combined state. All these compounds
are minerals, inorganic substances of definite composition found on
the earth. A mineral such as mercuric sulftde, HgS, or rock salt,
NaCl, from which an element, usually a metal, may be profitably
extracted is an ore.
How metals are extracted from ores. In mining operations, a
mineral is first separated mechanically from the rock, or gangue
(gang) , with which it is mixed by a process known as ore-dressing.
The particular process used depends upon the differences in prop-
erties between the ore and the gangue. Low-grade ores must be con-
centrated to reduce the cost of extracting the metal. Various methods
of ore-dressing, such as hand or gravity separation, leaching, mag-
netic, and flotation, are described in the discussion of specific ores.
The process of extracting a metal from its ore is called metallurgy.
In general, the metallurgy of any metal depends upon the nature and
purity of the ore, the properties of the metal to be extracted, and
the cost of the processes. The four most widely used metallurgical
processes are: (1) electrolysis; (2) reduction by means of carbon;
(3) roasting, that is, heating the ore to change it to its oxide by
oxidation (usually roasting is followed by reduction with carbon) ;
(4) aluminothermy, that is, reduction with aluminum powder.
The most widely used metal obtained by electrolysis is alumi-
num. Iron is the chief metal obtained by reduction, and copper and
zinc are the most common metals obtained "by roasting followed by
reduction. Chromium and manganese are made by aluminothermy.
These four metallurgical processes are discussed in Connection with
aluminum, iron, copper, zinc, and chromium.
370
METALS AND THEIR CHEMICAL ACTIVITY
371
Characteristic physical properties of a metal. Metals have many
physical properties in common. A brief study of these characteristic
properties helps us to realize why the classification of elements into
metals and nonmetals has aided the development of chemistry.
1) A metal is lustrous. That is, a metal has a definite shine. The
luster of gold, silver, nickel, and copper is well known.
2) A metal is malleable. That is, a metal can be hammered into
thin sheets such as gold leaf and tin foil. The most malleable metal
is gold, which has been beaten into sheets so thin that 300,000 of
them placed one above the other make a pile only one inch thick.
Carbon, a nonmetal, breaks when hammered; it is brittle.
3) A metal is ductile. That is, a metal can be drawn into wire.
Platinum, for example, can be drawn into such a fine wire that it
cannot be seen by the unaided eye. The degree of ductility of metals
varies greatly. Sulfur, a nonmetal, cannot be drawn into wire.
4) A metal is a good conductor of heat. Place one end of a copper
wire a foot long in a flame and notice how quickly your fingers hold-
ing the other end of the wire become warm.
5) A metal is a good conductor of electricity. Although silver is
the best conductor of electricity, copper is used most generally be-
cause of its much lower cost. On the basis of evidence now available,
scientists think that good conductors of electricity, including metals
in general and copper and silver in particular, are composed of
atoms whose outer rings have only a few electrons loosely held. These
(left) A prospector in Utah examines a sample of uranium ore. (right) Testing the
ore with acid for the presence of copper and limestone.
I'hotogmph*, Standard Oil Company (Ar./.)
Pouring molten copper from an electric
furnace into a mold. The "melt" con-
sists of both virgin metal and high-
grade scrap.
/>'< xeurch Association
electrons can, therefore, move along to the next atom, and so on,
producing a flow of electricity. Nonmetals, such as sulfur, have in
their outer rings a large number of electrons, not so free to move.
Hence nonmetals are poor conductors of electricity.
6) Other properties of a metal. All metals, with the exception of
mercury, are solids at ordinary temperatures. They range in melting
points from mercury ( — 39° C.) to wolfram (3380°C.) . Metals dif-
fer widely in tensile strength, that is, the ease with which they can
be pulled in two. Some steel has a tensile strength of 500,000 Ibs.
per sq. in. cross section.
They range in density from lithium (a little more than half as
heavy as water) , whose density is 0.53, to osmium, whose density is
22.5. Most metals are gray in color, the two most common excep-
tions being copper (red) and gold (yellow) . They range in hard-
ness from lithium, which is as soft as wax, to others which are very
hard. All metals are crystalline.
Carbide and Carbon Chemical* Company
A flotation cell in which the ores
of non-ferrous metals are con-
centrated. Unwanted impurities
settle to the bottom of the cell
and the ore is skimmed off with
the froth.
372
METALS AND THEIR CHEMICAL ACTIVITY
373
Characteristic chemical properties of a metal. Metals also possess
many chemical properties in common.
1) Certain metals unite with oxygen, forming oxides that are basic
anhydrides. The burning of magnesium, for example, produces mag-
nesium oxide, which is the anhydride of magnesium hydroxide.
2Mg -f O> -» 2M gO; MgO + H2O -> Mg(OH)2
2) Certain metals unite with water, forming either bases or oxides,
and liberating hydrogen. Thus steam passed over hot iron forms iron
oxide arid hydrogen, whereas calcium reacts with water, forming
calcium hydroxide and hydrogen.
3Fe + 4H2O
Ca + 2H2O
> Fe304 4- 4H2 T
> Ca(OH)2 + H2
3) Certain metals decompose acids, liberating hydrogen or other
gas.
Zn + 2HC1
Cu + 2H2SO4
3Cu + 8HNO3
CuS04 + 2H20 -f SO, t
> 3Cu(NOs)2 4- 4H20 -f 2NO
4) Metals combine with nonmetals, forming salts. For example,
they unite with chlorine, sulfur, and bromine, forming chlorides,
sulfidcs, and bromides, respectively.
5) Certain metals unite, with bases, liberating hydrogen. Thus,
both aluminum and zinc react with sodium hydroxide, liberating
hydrogen.
2A1 + 6NaOH -* 2Na,<AlO3 (sodium aluminate) -f 3H2 T
Zn -f 2NaOH -> Na2 ZnO2 (sodium zincate) + H2 \
Copper and Brass Research Association
These plates of impure
copper are about to be sus-
pended in a tank contain-
ing copper sulfate and
sulfuric acid. The copper
will be refined by an elec-
trolytic process described
in Chapter 28.
374 NEW WORLD OF CHEMISTRY
Metals may be listed in the order of their replacement power.
It is generally known that gold does not tarnish in air and is not
acted upon by any one acid, It is equally well known that iron,
on exposure to air, is oxidized readily. We could, if it served any
useful purpose, arrange all the metals in the order of their ability
to resist oxidation.
A more useful arrangement is based upon the ability of one
metal to replace another from a solution of a salt of the latter. For
example, when an iron nail is placed in a solution of copper sulfate,
the iron becomes coated with a layer of pure copper, and iron sul-
fate is formed.
Fe + CuSO4 -» Cu + FeSO4
Similarly, copper placed in a solution of silver nitrate becomes
coated with pure silver and will, itself, go into solution as copper
nitrate. The reverse reactions will not take place under normal con-
ditions. That is, a copper nail placed in an iron sulfate solution
will not deposit iron.
If we try these experiments, using a number of different metals
and their salts, we can arrange the common metals in a definite
replacement series. This replacement series of the common metals is
also called the electrochemical series, and the electromotive series.
How the electron theory explains the replacement series. Metals
differ in their tendency to lose electrons and become ions. When
free iron (Fe°) replaces copper from a copper sulfate solution, the
following electron reaction occurs:
loses 2 electrons
Fe° + Cu++(SO4)~ -+ Fe++(SO4)-- + Cu°
s, ^
gains 2 electrons
Electrically neutral iron loses two electrons and becomes ionic
Fe++, which is positively charged. Ionic copper from the copper sulfate
solution gains two electrons, becomes electrically neutral, and precipi-
tates out as free copper atoms. The sulfate ions remain unchanged.
Potassium
Sodium
Calcium
Magnesium
Aluminum
Zinc ^
Chromium
Iron
REPLACEMENT SERIES! NkTien
OF THE | lead
COMMON METALS 1 ^I0*"1
Copper
Mercury
Silver
Platinum
Gold
.aluminum
copper
,zmc
copper
zinc
zinc
Fead
Fig. 85. A strip of metal, placed in a solution of a salt of a metal below it in
the replacement series, replaces the less active metal, which precipitates onto the
strip.
The only change that takes place, then, is a transfer of two electrons
from free iron to ionic copper. Thus it seems that iron has a greater
tendency than copper to lose electrons. That is, iron is more metallic
than copper, and hence appears higher up in the replacement series.
If we place iron in a solution of calcium chloride, no reaction takes
place because the tendency of iron to lose electrons is less than that
of calcium.
Fe° + Ca++ + 2C1~ -* no reaction
The elements high up in the replacement series are so typically
metallic and have such a great tendency to lose electrons, that even
light causes them to throw off electrons. This fact is made use of in
the photoelectric cell. Such a cell is frequently lined with a thin film
of potassium, rubidium, or cesium, the most active metal known.
When light strikes this film, it throws off electrons, which travel to
a positively charged plate in the center of the cell. A very feeble
electric current is produced. This feeble current, whose strength
depends upon the intensity of the light that strikes the cell, may be
amplified and thus made to control larger supplies of energy. This
amplified current may close a switch, called a relay, which will start
a motor and open a door, or count people going through a passage.
The photoelectric cell is a vital part of sound-motion-picture and
television equipment, and it is used also in the transmission of pic-
tures by wire. Photoelectric cells are used in controlling the "blow"
of Bessemer converters in making steel.
Selenium, an element belonging to the sulfur family, was used in
certain "electric eyes." Selenium is a good electric insulator in the
dark, but in light it conducts an electric current to some extent.
Later it was replaced by copper covered by a thin film of Cu,O.
vacuated bulb
anode
cesium film
Fig. 86. Construction
of a photoelectric cell.
376 NEW WORLD OF CHEMISTRY
The replacement series of the nonmetals. In studying the halo-
gens, we learned that they, too, could be arranged according to
their ability to replace one another from solutions of their salts.
Thus, when chlorine is added to a solution of sodium bromide,
bromine is liberated and sodium chloride is formed. According to
the electron theory, this reaction is explained as follows:
gains 1 electron
SNa+Cl" + Br2°
loses 1 electron
Chlorine is more typically a nonmetal than is bromine. There-
fore, it has a greater tendency to gain electrons. Free chlorine, which
is neutral, takes one electron from the bromine ion and changes to
Cl~, that is, it goes into solution. The bromine ion, on the other
hand, after losing one electron, becomes electrically neutral, changes
to the atomic form, joins with another atom of bromine, and is liber-
ated as a free bromine molecule. Other nonmetals also may be
grouped in a series according to their replacement powers.
The replacement series of the metals is a useful tool. Under-
standing the replacement series of the metals is of great value in
studying these elements. Those above iron in the series, are so very
active that they are never found free, while those below iron occur
both in the free state and chemically combined. As we go down the
series, the tendency of the metal to lose electrons diminishes, and
hence the tendency to oxidize and to react with water or acids
diminishes also. Thus gold and platinum, which are at the bottom
of the list, do not oxidize in air even when hot, and are not attacked
by water or even by any one acid. Therefore, knowing the position
of a metal in this series, you can predict fairly well its chemical prop-
erties.
Why hydrogen is included in the replacement series of the metals.
Hydrogen, which from its physical properties could never be con-
sidered a metal, belongs in the replacement series because its ion is
usually positively charged and behaves chemically as a metal. All
metals above hydrogen in the replacement series liberate hydrogen
from acids. Those metals below hydrogen require oxidizing acids
to dissolve them and liberate gases other than hydrogen from these
oxidizing acids. For example, when copper reacts with nitric acid,
nitrogen dioxide is given off.
Humphry Davy isolates potassium and sodium. At the head of
the replacement list of the metals are the alkali metals, lithium,
Sir Humphry Davy (1778-1829), the son of
a poor wood carver, was born at Penzance,
Cornwall. The poet, Samuel Coleridge, de-
clared that if Davy "had not been the first
chemist, he would have been the first poet
of his age."
National Portrait Gallery, London
potassium, and sodium. Two of these elements had been known for
a long time as part of the alkaline compounds, potash and soda, be-
fore they were isolated in a pure state. The alkali metals were orig-
inally found in the ashes of certain plants. The name was taken from
the Arabic al, meaning the, and quili, meaning ashes.
Because of the extreme activity of these metals, many unsuccessful
attempts had been made to isolate them. After the discovery, in 1800,
of the galvanic current and the invention of the electric battery
which soon followed, scientists used this new force in an effort to
isolate sodium and potassium.
Humphry Davy, an eminent English chemist, rose from humble
beginnings to knighthood. He was the first to prove chlorine an
element and, incidentally, he was versatile enough to spend his lei-
sure hours writing fairly good poetry. In 1807 he sent the energy of
150 electric cells through molten potassium hydroxide. At the nega-
tive platinum electrode, Davy saw globules of a silvery substance
form, and then spontaneously catch (ire. "His joy knew no bounds,
he began to dance, and it was some time before he could control
himself to continue his experiments."
London received Davy's isolation of potassium as another wonder
o£ the world, and he was lionized. Some people paid 100 dollars to
attend his lectures on chemistry. Soon afterward, Davy obtained free
sodium in the same way, and lithium from fused LiCl.
How sodium is prepared for industrial use. The most recent
method of obtaining sodium in large amounts differs somewhat from
the method originally used by Davy. An electric current is sent
through melted sodium chloride in a cell, such as is shown in Fig.
377
sodium collect!
here
melted
sodium
rises
to DC source
chlorine
melted NaCI
metal screen
MM head
Fig. 87. Downs cell for
the preparation of sodium.
The cell was invented in
1924 by J. C. Downs, an
American chemist.
iron or copper cathode
«>$* graphite anode
87. Sodium ions travel to the iron or copper cathode, gain electrons,
become sodium atoms, and collect as a mass of metallic sodium,
which is drawn off from time to time. Chlorine ions travel to the
graphite anode, lose electrons, and become gaseous chlorine, which
leaves the apparatus as shown. This entire process is continuous. Po-
tassium and lithium may be prepared in this same way using melted
potassium chloride and lithium chloride.
Physical properties of potassium and sodium. Sodium is a soft,
silvery metal that melts just below the boiling point of water. Po-
tassium, which is also soft and silvery white, melts at an even lower
temperature. They are both lighter than water. Strangely enough,
these two solids when mixed form a liquid alloy at ordinary tem-
peratures.
Some chemical properties of sodium and potassium. The electron
pictures of lithium, sodium, and potassium are shown in Fig. 88.
From these pictures we can tell that the valence of each of these
elements is one. Each reacts with nonmetals, forming salts, and, on
exposure to the oxygen of the air, each is quickly tarnished with a
coating of its oxide.
Since sodium has only one electron to lend, and oxygen must
borrow two electrons to complete its outer ring, two sodium atoms
combine with one atom of oxygen, and the oxide of sodium is,
therefore, Na2O. Sodium peroxide, Na.X).,, (Na — O — O — Na) , is
formed when sodium is heated in air free from carbon dioxide.
Because of its extreme activity, sodium cannot be kept exposed to
air or under water. It is usually stored under kerosene, because kero-
sene contains no oxygen. At high temperatures both sodium and
lithium combine with hydrogen to form hydrides which react with
water liberating hydrogen.
LiH + H2O -» LiOH + H2 1
Fig. 88. Structure of the atoms of the alkali metals.
Lithium .'""'"^*
Sodium x C —
Potassium x^- "~
378
METALS AND THEIR CHEMICAL ACTIVITY
379
What has been said about the chemical activity of sodium applies
also to other alkali metals, namely, potassium, lithium, cesium, and
rubidium.
How the alkali metals are used. In a recent year 150,000 tons of
sodium metal were used in the manufacture of several compounds,
such as sodium peroxide, sodium cyanide, sodamide (NaNH,,) used
in making indigo, sodium hydride (NaH) used as a reducing agent
in removing surface oxides from steel, and several detergents. A
sodium-lead alloy is used in the manufacture of tetraethyl lead.
Liquid sodium, because it is an excellent heat conductor, is em-
ployed as a coolant in some nuclear reactors. Sodium is also used in
the hot cathode sodium vapor lamp, which gives twice as much
light as the common filament electric lamp using the same amount
of current. This type of lamp is used chiefly in outdoor lighting.
How can we test for the ions of sodium, potassium, and lithium?
If a clean platinum wire is dipped into a salt of potassium and then
placed in a nonluminous bunsen (lame, the flame becomes violet in
color. The flame of all sodium salts is a distinct yellow; that of all
lithium salts is red; cesium gives a bright blue flame. Since the pres-
ence of even a trace of a sodium salt will obscure the violet color
of potassium, the flame of a potassium salt can frequently be de-
tected only when viewed through a piece of blue cobalt glass, which
absorbs yellow light rays.
(left) Removing a sodium brick from a shipping drum, (right) Making sodium
pellets for laboratory use. Note the protective equipment in each case.
Ethvl Corporation
The bright line spectrum of lubricating oil (left) and of low-grade coal.
Sucli flame tests are used on other metals besides the alkali metals.
Thus, heated copper imparts a green color to the flame, and calcium
gives the flame an orange-red coloration.
How is a spectroscope used? In 1854, David Alter, a Pennsylvania
physician, described a method of detecting an element by the color
that it imparts to a flame. He also predicted the use of this method
in determining the presence of elements in the sun.
Five years later, Bunsen and Kirchhoff (kirK/hof) devised an in-
strument called the spectroscope, which has since become a very
powerful tool in the hands of chemists, physicists, and astronomers.
In I860, with the aid of this instrument, two new elements, cesium
and rubidium, were detected by Bunsen in a few grams of salt ob-
tained by the evaporation of 40 tons of spring water.
Nine years later, the element helium was discovered with the aid
of the spectroscope by Janssen and Lockyer independently. Helium
was found not on the earth, but in the sun, more than 90 million
miles away. Before the close of the century, this new element was
found on our own planet by Ramsay.
With the aid of the spectroscope, other elements, present in such
minute quantities that they had heretofore escaped discovery by
even the most delicate instruments of science, were finally brought
to light. Today, the spectroscope is used also in the study of the
complex structure of the atom.
This spectfOf rctpft makes Him records of spec-
tra, simitar to shown «bove» From tuch
tfMKtra, the composition of compounds may
An analytical chemist using elec-
tron diffraction equipment to
identify crystalline material such
as nickel oxide. Beams of elec-
trons aimed at the sample break
into a pattern which is made
visible by the equipment.
The principle of the spectroscope. To understand the principle
of the spectroscope, we must understand a few things about light.
Isaac Newton, in 1672, performed a classic experiment. He let a
beam of sunlight pass through a narrow slit into a dark room and
placed a glass prism in its path. A band of colors called the spectrum
was formed. This can be explained by remembering that light is a
form of energy which is transmitted by waves.
Sunlight is made up of light of various colors. Each color has a
different wave-length. Red has the longest wave-length (0.0000(i8
cm.) , and violet has the shortest wave-length (0.000040 cm.) . The
glass prism bent and split up, or refracted, the sunlight. The light
which was refracted, or bent, least was the red, and that which was
refracted most was the violet. Study the spectrum shown in the illus-
tration in color following page 382.
Though some self-luminous sources, such as the sun, have a con-
tinuous spectrum, as shown in the illustration, an incandescent va-
por or gas, such as heated sodium vapor or electrified neon gas, has
a discontinuous, or bright-line) spectrum. The glowing vapor of each
element has its own characteristic colored band of light. Thus, so-
dium vapor has one bright yellow line, lithium has one red and one
yellow line, and the vapor of iron has several hundred lines.
The spectroscope, the most essential part of which is a glass prism,
makes possible the quick analysis of incandescent vapors, and the
detection of the smallest trace of an element. Less than a millionth
381
382 NEW WORLD OF CHEMISTRY
of a milligram of sodium, and minute traces of poisons in blood
can be detected by spectroscopic analysis. A spectroscope is thus a
most useful tool to the chemists who specialize in the analysis of
many kinds of substances. Such specialists are called analytical chem-
ists.
Spectroscopic work of great precision is carried on today by means
of the spectrograph. This instrument differs from an ordinary spec-
troscope in that the observing telescope is replaced by a camera,
which makes a photographic record of the spectrum under examina-
tion. This permanent photographic record makes possible a more
careful analysis of the spectrum. Spectographic analysis is used in
steel and other alloys and compares favorably with the usual routine
quantitative analysis. Analysis of matter by the study of spectra is
called spectrum analysis. Because chemical manipulations are un-
necessary, all measurements may be made quickly.
YOU WILL ENJOY READING
Chemistry, Jan., 1945, pp. 37-43, and May, 1952. Published
by Science Service, Washington, D.C. Contain the original
papers describing the discovery of lithium, sodium, potassium,
cesium, and rubidium, and more information on the other
elements.
Mills, John. Electronics: Today and Tomorrow, pp. 84-94.
D. Van Nostrand Co., New York, 1944. Tells the story of the
photoelectric cell.
Pough, Frederick H. A Field Guide to Rocks and Minerals.
Houghton Mifflin Co., Boston, 1953. An excellent book for the
boy or girl interested in mineral-collecting as a hobby. Attrac-
tively illustrated.
USEFUL IDEAS DEVELOPED
1. A metal is an element that lends electrons. It has a posi-
tive valence, and its oxide is a basic anhydride.
2. An ore is a mineral from which an element may be
profitably extracted.
3. Ore-dressing is the separation of a mineral from the
valueless rock, or gangue, with which it is mixed.
4. Metallurgy is the process of extracting a metal from its
ore. The four chief metallurgical processes are (1) electrolysis,
(2) reduction by means of carbon, (3) roasting, or heating,
METALS AND THEIR CHEMICAL ACTIVITY 383
the ore to change it to its oxide, and (4) aluminothermy,
using aluminum powder.
5. Metals have certain characteristic physical properties.
A metal is (1) lustrous, (2) malleable, (3) ductile, (4) a good
conductor of heat, (5) a good conductor of electricity, (6) a
solid — mercury is an exception, (7) crystalline, and (8) vari-
able tensile strength.
6. Metals have certain characteristic chemical properties.
Some (1) unite with oyxgen, forming oxides which are basic
anhydrides, (2) unite with water, forming either bases or
oxides with the liberation of hydrogen, (3) decompose acids,
liberating either hydrogen or some other gas, (4) combine
with nonmetals, forming salts, and (5) unite with bases, lib-
erating hydrogen.
7. Metals may be so arranged in a replacement series that
each metal in the list will replace each metal below it from a
solution of its salt. According to the electron theory, this be-
havior is explained by the fact that the metals at the top of
the list lose electrons more easily than those at the bottom,
and hence go into solution more readily. The metals lower
on the list take these lost electrons, become electrically neutral,
and precipitate out as free metals.
8. Some nonmetals, such as the halogens, may also be ar-
ranged in a replacement series. According to the electron the-
ory, chlorine has a greater tendency to borrow electrons than
the other halogens and, hence, borrows electrons from those
below it in the series.
9. The flame test for identifying a metal consists of heating
a metal or one of its salts in a flame and noticing the color that
it imparts to the flame
10. The spectroscope is an instrument devised by Bunsen
and Kirchhoff. It is used to detect minute traces of elements
in incandescent vapors. The most essential part of a spectro-
scope is a glass prism, which disperses or breaks up the light
into colored lines which are characteristic for each element.
11. Spectrum analysis has been used in detecting the pres-
ence of rare elements, such as the inert gases of the atmosphere.
It is also used in the study of the structure of the atom.
USING WHAT YOU HAVE LEARNED
Group A
1. According to the electron theory, how do metals differ
from nonmetals?
384 NEW WORLD OF CHEMISTRY
2. What are four typical nonmetals and four typical
metals?
3. (a) Name a few metals that occur free, (b) Compare
their chemical activity with the chemical activity of metals
found combined.
4. (a) In what way does a mineral differ from an ore?
(b) Give an example of each, (c) Do minerals ever become
ores? (d) Illustrate.
5. (a) What are the four most general methods used in
metallurgy? (b) Give an example of an element extracted by
each of these methods.
6. What are the characteristic physical properties of metals?
7. (a) Arrange these eight metals in a replacement series:
Zn, H, Cu, Na, K, Pt, Au, and Pb. (b) On your list, check
those that will replace the hydrogen of dilute HC1.
8. Arrange the number of each metal opposite the letter
of the property of which it is an outstanding example.
a) Best conductor of electricity 1) Cu
b) Extremely ductile 2) Au
c) Most malleable of all 3) Li
d) Lightest metal 4) Hg
e) Heaviest metal 5) W
f) Liquid metal 6) Os
g) Extremely high melting point 7) Ag
h) Reddish luster 8) Pt
9) Na
9. (a) Write five equations illustrating five chemical prop-
erties of metals, (b) In each case state the property.
10. According to the electron theory, why does an iron nail
become coated with Cu when placed in a solution of CuSO4?
11. In which of the following would a replacement reaction
take place? Complete the equations for such replacements.
a) Zn + Hg(NO3)2-* e) Cu + AgNO3 -»
b) Zn + CuSO4-> f) Cu + ZnSO4->
c) Zn + AgNO3 -» g) Cu + Hg(NO3)2 ->
d) Zn -f Pb(NO3)2 -» , h) Cu + Pb(NO3)a -»
12. Arrange a replacement sefies of some nonmetals you
have studied.
13. According to the electron theory, explain why Br2 lib-
erates free I9 from a solution of KI.
METALS AND THEIR CHEMICAL ACTIVITY 385
14. An element X is not found free. It attacks warm water,
liberating H2, and tarnishes readily in air. Where would you
place it in the replacement list of metals? See Tat>le 12.
15. (a) What are two elements other than Na that belong
to the sodium family? (b) Why are they called alkali metals?
16. (a) How and by whom were Na and K first isolated?
(b) How was the news of this achievement received?
17. (a) What are three ways in which Na is similar to K?
(b) one way in which it is different?
18. By a diagram, show the present method of obtaining
free Na.
19. Make a diagram of the atom of Li, and, from this
diagram, state its (a) valence, (b) chemical activity, and
(c) atomic number.
20. Using an equation, describe the action of Na on H2O.
21. What weight of NaOH must be decomposed to produce
69 g. of pure Na?
22. A piece of Na is placed in H2O, and 336 ml. of H2 are
collected. What weight of Na took part in this chemical
change?
23. Which will require more H0O in dissolving completely
45 g. of Na or 79 g. of K?
24. Copy and complete. Do not write in this book. Na is
stored under .... When Na is exposed to air, the formula of
the compound formed is .... Na is used as a catalyst in the
preparation of .... An instrument that makes use of the ease
with which Cs loses electrons is the .... The color imparted
to a flame by K vapor is . . ., Li a ... color, and Na a ...
color.
25. (a) List four uses of sodium, and (b) three uses of
other alkali metals.
26. How is the spectroscope used in astronomy?
27. Name three elements discovered by means of the spec-
troscope.
28. How can the spectrum of sunlight be obtained?
386 NEW WORLD OF CHEMISTRY
29. What is the difference in the appearance of the spectrum
of a luminous solid and the spectrum of an incandescent
vapor?
30. Hydrogen is included in the replacement series of the
metals. Why?
Group B
31. Zn appears higher than Fe in the replacement series.
What is the reason for coating Fe with Zn to prevent corrosion?
32. Devise an experiment for obtaining Cu from CuSO4
solution.
33. Na and K kept under kerosene for some time lose their
silvery luster. Explain.
34. Discuss the use of spectroscopy in crime detection.
PROJECTS, ACTIVITIES, AND INVESTIGATIONS
1. Purchase or borrow a photoelectric cell and connect it
in such a way that, when a flashlight is directed against it, a
bell will ring. Explain its action to the class. What use would
you make of a photoelectric cell?
2. Make your own replacement series of some or all of the
following metals: aluminum, copper, chromium, lead, tin,
zinc, and calcium (if you can obtain a small piece from your
teacher) . Use whatever suitable chemicals you can find around
the house such as vinegar. Report your results.
3. Prepare an illustrated ten-minute talk on mineral-col-
lecting as a hobby. Bring some of your specimens to class.
4. Write a 300-500 word essay on one of the various ana-
lytical methods used by chemists today such as (a) wet
method, (b) spectrophotometry, (c) chromatography, (d) tracer
technique.
ALUMINUM:
MOST COMMON
OF LIGHT METALS
. . . / believe I speak for the vast ma-
jority of all scientific men. Our ob-
ject is riot to make fobs and divi-
dends. These are a means to an end,
merely incidental. We wish to abol-
ish drudgery, discomfort, and want
/row the lives of men, and bring
them pleasure, comfort, leisure and
beauty. Harold C. Urey, 1934
The world receives a valuable gift. In 1825, Hans Christian
Oersted (lir'stetfe) , a Danish scientist, announced that he had isolated
aluminum by gently heating aluminum chloride and potassium
amalgam. In 1827 Woehler repeated Oersted's experiments without
success. Woehler finally obtained aluminum by substituting metallic
potassium for the potassium amalgam used by Oersted.
One of the first of Woehler's American students in Germany was
Professor Jewett, of Oberlin College. He brought back to America
the story of Woehler's isolation of that extremely light, silvery metal,
aluminum, fewett was fond of talking to his classes of this strange
metal, which no one had as yet been able to obtain inexpensively in
spite of its great abundance in minerals.
One day, as Professor Jewett spoke of the fortune that awaited the
man who was able to develop a simple method for extracting alu-
minum, one of the students nudged his young classmate, Charles
Martin Hall. Chemistry had captivated Hall, and his classmates
had known him to make all sorts of experiments, hoping to make a
great discovery some day. Here was his chance. His response to that
nudge was, "I am going after that metal," and Hall went to work at
once in his father's woodshed
387
388 NEW WORLD OF CHEMISTRY
Hall attacked his problem scientifically. He knew that only the
most active metals, such as sodium and potassium, were reducing
agents strong enough to liberate aluminum from aluminum chloride.
A1C13 + 3K -> 3KC1 + Al
Potassium had been used in the method developed by Oersted
arid used later by Woehler, and sodium had been substituted tor po-
tassium by the French chemist, Henri Sainte-Claire Deville (saNt-
klar' de-vel') . But both potassium and sodium were too expensive
to use in a commercial method. Hall finally discarded all methods
that depended upon the action of a reducing agent and turned to
electrolysis, in spite of his knowledge that Davy, who had isolated
the alkali metals by electrolysis, had failed to get pure aluminum in
this way.
Aluminum oxide, called alumina, was the natural starting point.
Alumina in hydrated form is the chief component of bauxite, the
richest ore of aluminum. But alumina has an extremely high melting
point. To melt alumina was commercially impracticable. But if an
electric current was to liberate free aluminum from it, alumina had
to be either melted or dissolved. Perhaps (the thought came to Hall
in one of those flashes of genius) some mineral that would act as a
solvent for aluminum oxide might be found. After trying a number
of minerals, he came across a milky-white, glassy solid called cryolite.
He melted this with some difficulty and then threw in some alumina.
The alumina dissolved readily. He passed a current through the
solution of alumina in cryolite and, to his intense joy, found that
metallic aluminum was deposited at the cathode
On February 23, 1886, he burst excitedly into the laboratory of
Professor Jewett and, holding a few aluminum "buttons," exclaimed,
Alum
Charles Martin Hall, the discoverer of the
electrolytic process of producing alumi-
p, «o c ,- , + I. .1. 4. ^busbar
Fig. 89. Cross section of a '.. \m\ i,i D ..' ^,^Mt«» JL A! o
* , .. . . . . M TT T • rTx^'^crusF or AUw3
furnace in which aluminum is I] 1 _ | _ I ^Jr[
produced commercially by the ^j^jjjAjjjLj^^ AI2O3 In
Hall process. carbon anode + ^g^^^^^J^Tn fused^cryolite
carbon lining cathode aluminum
"I've got it!" Hall soon obtained a patent on his process and two
years later the Pittsburgh Reduction Company, which later became
the Aluminum Company of America, was founded. In 1914 Hall
died, world-famous and a multimillionaire. He left most of his for-
tune to Oberlin College and other educational institutions.
Discoveries often result from social needs. Hall was 22 when he
produced aluminum. Exactly two months after Hall had handed his
teacher the first samples of his aluminum, another chemist, Paul
Heroult (a-roolt') , also 22 years of age, applied for a patent in France
on an identical commercial method of preparing aluminum.
This is not a rare example of simultaneous discoveries in the his-
tory of science. Advances in science often are made in different parts
of the world at almost the same time. They are frequently the final
step in a long series of experiments conducted by many research
workers in widely separated laboratories. The scientist who is fortu-
nate enough to publish his discovery first is recognized as the hero
of a battle in which many other soldiers have been engaged. The
heroes of science, on the whole, concede this element of good for-
tune. Can you think of another such instance in the story of scientific
advance?
Metallurgy of aluminum by the Hall process. Hall's process is still
in use. The electric furnace used is an open cell about 25 to 30 feet
long, lined with carbon, which constitutes the cathode. Powdered
cryolite mined in Greenland or made synthetically is placed in the
cell, and as the current passes through it, it melts. Aluminium oxide,
or alumina, is a white powder obtained by refining bauxite ore. It
is added to the molten cryolite and immediately dissolves. The
aluminum oxide dissociates, forming positive aluminum ions and
negative oxygen ions.
A12O3 -» 2A1+++ + 3O—
Carbon rods are suspended in the molten aluminum oxide solution
and act as the anode. When the circuit is closed, the aluminum ions
travel to the cathode, where they obtain electrons which change
them from aluminum ions to free aluminum. This free molten alu-
minum then settles to the bottom of the cell. Later, a hole at the
389
390 NEW WORLD OF CHEMISTRY
bottom of the cell is unplugged, and the molten aluminum is tapped
off into large ladles and cast in molds, in which it solidifies as pig
aluminum. The oxygen ions, in the meantime, have traveled to the
anode, given up their electrons, and changed to free oxygen. This
oxygen combines with the carbon of the anode and forms carbon
dioxide, given off as a gas.
c -f o2 -» co2 T
The process is continuous. Aluminum oxide is added, aluminum is
removed, and the carbon anodes are replaced from time to time.
The original cryolite, Na3AlF«, though it contains aluminum, does
not decompose. It acts only as a solvent. Many of these electrolytic
cells are joined in series. Commercial cells produce about 500 pounds
of 99+ percent pure aluminum per day.
Bauxite, found in large amounts in Surinam (Dutch Guiana) ,
and British Guiana — Arkansas leads the United States in the pro-
duction of this ore — contains a fair percentage of the oxides of iron,
silicon, and titanium. If these impurities are not removed before
the bauxite is added in the electric furnace, the aluminum produced
is impure.
After the aluminum is drawn from the electrolytic cells, the pig
metal is remelted so that the nonmetallic impurities may be skimmed
off. If aluminum alloys, rather than pure aluminum, are desired, the
alloying may be done during the remelting. The chief alloying ele-
ments include copper, magnesium, manganese, silicon, zinc, iron,
nickel, and chromium.
The physical properties of aluminum. Aluminum is silvery white
in color and is one of the lightest of the common metals. It is only
one-third as heavy as iron. It is very malleable and ductile and com-
pares well with both silver and copper in the ease with which it con-
ducts both heat and electricity. It can be worked readily; that is, it
can be cast, rolled, forged, extruded, machined, or drawn. Parts can
be joined by welding, brazing, and riveting.
The chemical properties of aluminum. The atomic weight of alu-
mium is 27. Its atomic number is 13; hence it has only three elec-
trons in its outside ring. It is, therefore, a metal with a valence of
plus three. Aluminum is an amphoteric element and may act as either
an acid or a base. It is attacked by strong bases as follows:
2A1 + 2NaOH + 2H2O -» 3H2 1 + 2NaAlO2 (sodium aluminate)
This sodium aluminate is the salt of aluminic acid, H3A1O3. Be-
cause of the reaction between aluminum and the strong bases, or
ALUMINUM AND MAGNESIUM
391
substances with basic reactions, such as washing soda, care should be
taken not to heat such substances in aluminumware.
Aluminum is attacked by nearly all acids, forming aluminum salts.
2A1 + 6HC1 -» 3H2 1 + 2A1C13
The surface of aluminum oxidizes rapidly in air, forming alu-
minum oxide, AL,O3. This extremely thin, transparent, but tough
film acts as an excellent protective coating and, unlike iron rust,
adheres firmly to the surface of the metal, thus preventing further
oxidation unless the coating is perforated. Alclad is a sheet of alu-
minum alloy such as duralumin covered with a layer of aluminum.
It resists corrosion very well.
Tremendous growth of the aluminum industry. Before the Hall
process was introduced, aluminum was not used widely, because of
the great cost of preparing it. It is a far cry from the world produc-
tion of two tons of aluminum in 1859 at 17 dollars per pound to the
more than one million tons of this metal produced in the United
States and Canada alone in a recent year at about 20 cents a pound.
The Hall process gives primary aluminum, that is, metal pro-
duced directly from an ore or ores. But a significant and perhaps
increasing source of aluminum, and other metals as well, lies in
secondary sources,, that is, sources from which a metal is recoverable
from one use for reuse in another. Secondary or scrap aluminum is
a very large source of pure aluminum. Production of metals from
Surface mining of bauxite in Surinam. Over-lying earth is first removed. The ex-
posed ore is then loosened by blasting and loaded into the mine cars.
Aluminum Company of America
An "aluminum skyscraper" in Pitts-
burgh, Pa. This 30-story building is
constructed of aluminum panels
mounted on a steel framework. The
ceilings, wiring, ventilation ducts,
doors, hardware and most of the
plumbing are made of aluminum.
secondary sources is one phase of an intelligent metals conservation
program. Such a program can do much to conserve natural resources.
By far the largest users of aluminum are the transportation indus-
tries. Great quantities of the "metal with wings" are used in the
construction of airplanes, streetcars, railroad cars, locomotives, steam-
ships, motorships, automobiles, trucks, buses, bicycles, and motor-
cycles.
Through decreasing the weight of such carriers as airplanes, rail-
road cars, and trucks, payload and, thereby, revenue can be increased.
At the same time, if the reduction in weight applies only to the total
gross weight of the vehicle, as in an automobile or bus, much less
energy is required to attain and maintain speed, much less energy
is lost in stopping, and operating costs are thereby reduced.
Electricity and aluminum. The electric industries use thousands
of tons of aluminum yearly in lines for the transmission of electricity
over long distances. For this purpose, aluminum cable with a steel
reinforcing core is used instead of copper. Since aluminum is lighter
than copper, fewer towers are required to support the cables. More
than 1.5 million miles of aluminum transmission cables carry elec-
tricity to almost all parts of this nation. Aluminum is used also in the
production of parts for electric equipment of many kinds, such as
392
ALUMINUM AND MAGNESIUM 393
vacuum cleaners and various household appliances, and particu-
larly in parts for radios and other electronic equipment.
Aluminum in the kitchen. One use of aluminum goes back to
1890, when the first aluminum cooking utensils were produced.
Since aluminum is an excellent conductor of heat, and at the same
time is very light, aluminum cooking utensils are very popular.
Certain alkaline foods and waters heated in aluminum may pro-
duce a superficial discoloration, which is harmless and readily re-
moved by a mild abrasive cleaner. Do not permit all the cooking
liquid in a lightweight aluminum utensil to boil off, for if this
occurs, a hole may be "burned" in the bottom of the utensil as a
result of the relatively low melting point of aluminum.
Because strong alkalies attack the protective coating of alumi-
num oxide that forms on the surface of aluminum as well as the
aluminum itself, cooking utensils made of the metal should not be
scrubbed or polished with harsh alkali cleaners. To clean aluminum
utensils use soap and water or mild abrasive cleansers only.
Other uses of aluminum. In the packaging of foods and other com-
modities, aluminum foil has almost entirely replaced tin foil. Candy
bars, chewing gum, cream cheese, camera film, and countless other
articles go to market in shining dress. Aluminum foil coated with a
plastic film is suitable for the packaging of almost any kind of food.
Aluminum leaf is used in photoflash lamps.
Collapsible tubes made of aluminum carry such items as shaving
cream, tooth paste, and cosmetics, while vital serums and various
other pharmaceutical preparations are packed in glass bottles with
aluminum seals. Aluminum paints are widely used for protecting
both wood and metal. Aluminum foil is used for home insulation.
Aluminum Company of America
Circular aluminum blanks
being removed from the
conveyor of an annealing
furnace. These blanks will be
used in the manufacture of
cooking utensils.
394
NEW WORLD OF CHEMISTRY
ALUMINUM AND MAGNESIUM
395
Much furniture and many decorative articles for the household are
made of aluminum.
The buildings of Rockefeller Center in New York contain more
than 1000 tons of aluminum in the vertical panels between windows.
In the finishing of steel, large quantities of aluminum are used in
removing oxides from the molten steel. The largest use of aluminum
is in the form of alloys of much greater strength than pure aluminum.
Thermit is used in welding. Attempts to reduce aluminum oxide
with carbon failed because of the great attraction of aluminum for
oxygen. Aluminum is a powerful reducing agent, especially when
it is in the form of a fine powder. Because a powder has a much
greater reacting surface than a solid, a powder makes possible a more
intense chemical reaction than the same weight of the same solid in
larger pieces.
When a mixture of powdered aluminum and iron oxide is ignited
by means of a fuse, such as a strip of magnesium ribbon, a chemical
reaction takes place at once in which the aluminum takes the oxygen
away from the iron oxide, leaving a residue of pure iron.
8A1 + 3Fe3O4 -» 4A12O3 + 9Fe
The heat of this reaction is so great that the iron formed is molten.
This mixture of aluminum and iron oxide, known as thermit, is used
in welding broken propeller shafts, rudder frames, locomotive parts,
and in situations where repairs must be made on the spot. It is also
used in one type of incendiary bomb.
The metals chromium and manganese (and wolfram, vanadium,
molybdenum, silicon and boron) may be extracted from their oxides
or ores by aluminothermy, an aluminum reduction similar to the
thermit reaction. The equations for the reduction of chromium and
manganese ore are:
Cr2O3 + 2A1 -> A12O3 + 2Cr
3Mn3O4 + 8A1 -» 4A12O3 + 9Mn
This is a common way of manufacturing or producing such metals.
What are alums? When potassium sulfate is dissolved in a solution
of aluminum sulfate, the two salts combine and crystallize out as
potassium aluminum sulfate.
K2SO4 + A12(SO4)3 + 24H2O -> 2KA1(SO4)2 12H2O
This compound is called common alum. It is one of a group of salts,
called the alums, which resemble one another in the eight-sided form
of their crystals, in their solubility in water, and in their type for-
mula. The type formula of an alum is XY (SO4) L> • 12H2O. In it X may
be K, Na, or NH4, and Y may be a trivalent element, such as Al, Fe,
or Cr. An alum is a double salt, that is, a salt containing two metals
and one acid radical.
Common alum has a sweetish taste and is used in "Foamite," mor-
danting (see page 605) , and water purification. Alum is used also
in making alum baking powder and as an astringent, a substance that
contracts skin tissues. It is used in the sizing of paper. Chrome alum
is used in the tanning of leather, during which skins or hides are
made softer, and more resistant to the action of bacteria.
Aluminum hydroxide, its preparation and properties. When any
soluble aluminum salt is added to water or to ammonium hydroxide,
aluminum hydroxide, a white, gelatinous precipitate, may be formed.
The aluminum salt is said to hydrolyze, since, as in the case of the
solution of aluminum chloride in water, an acid and a base are
formed.
A12(SO4)3 + 6NH4OH •
A1C13 + 3H2O
- 2A1(OH)3 1 + 3(NH4)2SO4
• 3HC1 + A1(OH)8 1
ALUMINUM
Use by approximate
Industrial, agricultural,
and mining machinery
396 NEW WORLD OF CHEMISTRY
Aluminum hydroxide may act either as an acid or as a base. When
it comes in contact with a base, it acts as a weak acid and combines
with the base.
HaAlOs + NaOH -> NaAlO2 (sodium aluminate) + 2H2O
A1C18 is the active constituent, about 15 percent, of many body
deodorants and antiperspirants.
Aluminum sulfate in the purification of water. As you know,
aluminum sulfate, when it is added to water, hydrolyzes, forming
jelly-like aluminum hydroxide.
A12(S04)3 + 6H20 -» 2A1(OH)3 j + 3H2SO4
This aluminum hydroxide gradually settles to the bottom and carries
down with it any particles that are floating in the water, including
bacteria, industrial wastes, and fine clay. This process, used in the
purification of water, is called coagulation and is the first step in
clearing water of its turbidity. Coagulation does not remove dis-
solved impurities.
Ferrous sulfate is sometimes used instead of aluminum sulfate,
because it also forms a gelatinous precipitate, ferrous hydroxide, in
water. But ferrous hydroxide rapidly oxidizes to ferric hydroxide,
which is also gelatinous. Hence, it is the ferric hydroxide that actu-
ally reduces the turbidity, or clarifies the water.
FeS04 + 2H20 -> H2SO4 + Fe(OH)2 j (ferrous hydroxide)
Aluminum oxide — ore, gem, abrasive, refractory. When alumi-
num hydroxide is heated, it forms a white, insoluble oxide of alumi-
num, which melts above 3600° F.
2A1(OH)3 -> A12O3 + 3H2O
As hydrated oxides, A12O3 • 3H2O and A12O8 • H2O are found widely
distributed in the aluminum ore, bauxite. The precious gem stones
ruby and sapphire are composed of alumina, colored by the pres-
ence of small amounts of metal oxides. Successful methods have
been developed to prepare synthetic rubies and sapphires by melting
pure aluminum oxide in the heat of an oxy hydrogen flame. These
synthetic stones cannot easily be distinguished from natural gems.
The production of synthetic rubies and sapphires is increasingly
important, for they are used as bearings in watches, electric indica-
tors, sensitive electric relays, and in thousands of other kinds of pre-
cision instruments.
ALUMINUM AND MAGNESIUM 397
Emery is a natural aluminum oxide, which is extremely hard and
can be used as an abrasive for grinding, polishing,